iraiv 


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MEMCAL    .S€1HI©(Q)L 


California  College  of  Pharmac^ 


AN  ELEMENTARY  TREATISE  ON 

QUALITATIVE  CHEMICAL 

ANALYSIS 


BY 

J.  F.^aLERS 

Professou  of  Chemistry,  Mercer  University,  Georgia 


California  CoHege  of  Pharmacy 


REVISED  EDITION 


GTNN  AND  COMPANY 

BOSTON  •  NEW  YORK  •  CHICAGO  •  LONDON 


Copyright,  1900, 1909 
By  J.  F.  SELLERS 


Alili  RIGHTS  RESERVED 


CINN   Sz   COMPANY  •  PRO- 
PRIETORS .  BOSTON  .  U.S.A. 


3 


PREFACE 

So  many  books  on  analytical  chemistry  are  already  in 
print  that  the  question  may  be  raised  whether  it  is  wise 
to  add  still  another  to  their  number;  and  therefore  the 
author  desires  to  present  the  following  reasons  which  seem 
to  him  to  justify  the  publication  of  the  present  work. 

Most  writers  on  analytical  chemistry  have  gone  either 
to  the  one  or  the  other  of  two  extremes.  First,  there  are 
those  who,  like  Fresenius  or  Prescott  and  Johnson,  have 
endeavored  to  cover  the  entire  field  and  to  include  the 
whole  detail  of  analytical  chemistry.  Their  works  are 
indispensable  to  teachers  and  to  students  who  make 
chemistry  a  specialty  ;  but  for  beginners,  who  may  not 
give  more  than  one  year  of  eight  or  ten  hours  a  week  to 
the  subject,  they  are  fq,r  too  voluminous.  On  the  other 
hand,  there  are  those  whose  ardor  for  brevity  has  led  them 
to  the  other  extreme  of  condensing  their  material  into 
"  tables  "  and  "  schemes,"  —  by  which  means  they  have 
magnified  the  empirical  and  have  minimized  the  rational 
aspect  of  the  subject,  to  its  considerable  detriment  as  a 
factor  in  liberal  education. 

In  order  to  avoid  either  extreme  the  writer  presents 
this  elementary  treatise  having  these  features  :  — 

1.  A  course  short  enough  to  be  digested  during  the 
time   allotted  in   an  ordinary  college  curriculum,  but  at 

iii 

39971 


IV  PBEFA  CE 

the  same   time  intended   to   magnify   the   scientific   and 
pedagogical  nature  of  analytical  chemistry. 

2.  A  course  both  practical  and  progressive,  —  practical, 
in  that  the  student  can  master  the  methods  and  principles 
of  chemical  analysis,  and  become  a  practical  analyst ;  pro- 
gressive, in  that  the  chief  aim  of  the  book  is  to  prepare  the 
student  thoroughly  for  advanced  university  work. 

3.  A  course  selected  from  the  most  recent  and  approved 
methods  recorded  in  the  best  literature  and  verified  by 
actual  application  in  the  author's  laboratory.  Among 
some  of  the  improved  methods  are  mentioned  :  — 

(a)  Reddrop's  application  of  normal  solutions  to  quali- 
tative analysis.      Chemical  News,  May,  1890. 

(h)  Hofmann's  separation  of  arsenic,  antimony,  and 
tin,  by  modification  of  Marsh's  test.  Fresenius'  Quali- 
tative Analysis,  1897  edition,  p.  299. 

(c)  Parr's  separation  of  aluminum,  chromium,  and  iron, 
by  means  of  sodium  peroxide.  This  method  commends 
itself  for  its  accuracy,  its  briefness  and  simplicity,  and  its 
certainty  in  detecting  aluminum.  Other  methods  depend- 
ing on  sodium  hydroxide  are  defective,  in  that  the  reagent 
itself  generally  contains  aluminuni  salts  ;  sodium  peroxide, 
by  reason  of  its  manufacture,  does  not  contain  perceptible 
traces  of  such  salts.     Journ.  Amer.  Chem.  Soc.,  19,  p.  341. 

(d)  Fresenius  and  E-uppert's  separation  of  barium, 
strontium,  and  calcium,  by  means  of  the  differences  of 
solubility  of  their  nitrates  in  ether-alcohol.  Fres.  Qual. 
Anal.,  p.  160. 

(e)  Hager's  separation  of  chlorine,  bromine,  and  iodine, 
by  means  of  the  differences  of  solubility  of  their  silver  salts 
in  ammonium  "  sesqui "  carbonate.   Fres.  Qual.  Anal.,  p.  378. 


PREFACE  V 

4.  A  course  free,  as  is  thought  wise,  from  the  mechan- 
ical schemes  in  qualitative  analysis.  To  this  end,  many 
of  the  usual  tables  of  separation  are  omitted,  and  in  their 
place  some  suggestive  hints  are  given  after  the  list  of 
reactions  for  each  group.  Thus  the  student  is  expected 
and  encouraged  to  exercise  his  judgment  in  selecting 
methods  of  analysis. 

5.  A  course  conformable  to  the  modern  dissociation 
theory  of  solutions.  For  example,  why  is  the  activity 
of  certain  acids  modified  by  adding  the  salts  of  those 
acids  J  or,  more  specifically,  why  is  the  solvent  power 
of  acetic  acid  decreased  by  adding  some  sodium  acetate  ? 

6.  A  course  giving  more  than  ordinary  emphasis  to 
the  spectroscope.  Though  spectroscopy  is  not  chemical 
analysis,  it  possesses  superior  advantages  over  the  chem- 
ical methods  in  these  particulars  :  — 

(a)  Methods  of  greater  exactness  and  readiness  of 
execution. 

(b)  Methods  superior  for  the  preliminary  detection  of 
the  alkali  and  alkali-earth  metals.  This  is  important, 
especially  when  the  alkali-earth  metals  are  combined  with 
phosphoric,  oxalic,  and  hydrofluoric  acids. 

(c)  Methods  superior  for  detecting  certain  metals,  which, 
under  some  conditions,  are  evasive;  e.(/.,  aluminum,  man- 
ganese, and  magnesium. 

It  is  obvious  that  the  study  of  the  theory  of  solution 
and  of  spectroscopy  may  either  be  taken  up  in  the  order 
of  the  text  or  reserved  for  the  last  work  in  the  course; 
and  also  that  these  subjects  may  be  omitted  entirely  if 
a  very  elementary  course  is  desired.  In  the  latter  case 
it  would  be  possible  also  to  omit  the  discussion  of  the 


VI  PREFACE 

analysis  for  the  metals  of  the  third  group  in  the  pres- 
ence of  phosphoric  acid,  and  those  portions  of  Part  II 
which  are  printed  in  small  type. 

The  discussion  of  solutions  in  the  brief  space  available 
in  this  book  is  necessarily  much  condensed,  and  possibly  it 
is  somewhat  abstract  and  uninviting  ;  but  in  the  author's 
opinion  its  introduction  is  desirable.  Its  purpose  is  to 
provide  the  student  of  qualitative  analysis  with  the  means 
for  a  rational  interpretation  of  many  apparently  irrational 
reactions,  and  to  help  prepare  him  for  the  next  stage  of 
his  chemical  education,  —  namely,  the  study  of  quantitative 
analysis,  —  where  the  application  of  the  laws  of  solutions 
is  more  abundant.  No  other  text-book  on  qualitative 
analysis,  within  the  author's  knowledge,  incorporates 
this  dissociation  theory  of  solution  ;  but  its  adaptability 
to  qualitative  instruction  is  shown  by  the  fact  that  during 
the  past  half  decade  many  teachers  of  the  subject  have 
devoted  more  or  less  time  in  their  lectures  to  the  practical 
application  of  the  theory. 

In  the  preparation  of  this  book  the  following  literature 
has  been  consulted  :  — 

1.  Many  of  the  smaller  text-books  on  qualitative  analysis, 
including  Noyes's,  Newth's,  and  Volhard  and  Zimmer- 
mann's. 

2.  Standard  works  on  general  and  analytical  chemistry, 
including  Watt's  Chemical  Dictionary ;  Eoscoe  and  Schor- 
lemmer's  Treatise  on  Chemistry  ;  Mendeleeff's  Principles  of 
Chemistry  ;  Ostwald's  and  Nernst's  works  on  physical  chem- 
istry ;  Vogel's,  Landaur's,  and  Eoscoe's  works  on  spectrum 
analysis  ;  Fresenius'  works  —  the  latest  editions. 

3.  Memoirs  in  American  and  foreign  chemical  journals. 


PREFACE  VU 

Grateful  acknowledgment  is  made  to  Dr.  E.  W.  Jones 
of  the  University  of  Mississippi,  for  his  painstaking  criti- 
cism of  the  manuscript  of  this  little  book.  The  author 
learned  the  chemical  alphabet  and  received  much  inspira- 
tion and  encouragement  from  this  excellent  teacher. 

Appreciative  mention  also  is  made  of  the  following  gen- 
tlemen :  Dr.  H.  C.  White  of  the  University  of  Georgia, 
for  valuable  suggestions  as  regards  the  adaptability  of  the 
book  to  elementary  college  work ;  Dr.  J.  W.  Mallet  of 
the  University  of  Virginia,  Dr.  J.  Stieglitz  of  the  Uni- 
versity of  Chicago,  and  Dr.  E.  Eenouf  of  Johns  Hopkins 
University,  for  opinions  concerning  modern  theories  of 
solution;  and  Mr.  H.  V.  Jackson  of  Mercer  University, 
for  general  assistance.  t  t^  a 

Macon,  Ga.,  September,  1900 

PREFACE    TO    SECOND    EDITION 

The  more  important  modifications  made  in  this  edition 
of  the  book  are  the  appending  of  13  pages  of  reference 
notes  (see  p.  163)  and  tables,  revision  of  several  of  the 
processes  of  separation,  and  correction  of  a  number  of 
typographical  errors. 

For  criticism  and  proof  reading  the  author  is  indebted 
to  many  of  his  friends,  among  whom  may  be  mentioned 
Professor  W.  H.  Emerson  and  Dr.  G.  H.  Boggs  of  the 
Georgia  School  of  Technology,  Dr.  J.  P.  Montgomery  of 
the  Mississippi  Agricultural  and  Mechanical  College,  Dr. 
Homer  V.  Black  of  the  University  of  Georgia,  Professor 
C.  W.  Steed  of  Mercer  University,  Professor  G.  P.  Shingler 
of  Emory  College,  and  Professor  Alexander  Smith  of  the 

University  of  Chicago. 

J.  F.  S. 
Macon,  Ga.,  May,  1909 


CONTENTS 


PART    I  — ANALYTICAL   OPERATIONS 

CHAPTEU 

I.    Introduction 


II.  Theory  of  Analytical  Operations    . 

III.  Methods  of  Analytical  Separation  . 

IV.  Flame  Coloration  and  Spectroscopy 
V.  List  and  Preparation  of  Reagents  . 

VI.  Systems  of  Analytical  Examination 


PAGB5 

1 

5 
26 
50 
65 
73 


PART    II  — REACTIONS    AND    SEPARATIONS 

VII.     Metals  of  Group  I         77 

VIIL     Metals  of  Group  II        82 

IX.     Metals  of  Group  III 100 

X.     Metals  of  Group  IV 114 

XI.    Metals  of  Group  V        121 

XII.    Metals  of  Group  VI 126 

Xiri.     Acids  of  Group  I 130 

XIV.     Acids  of  Group  II 141 

XV.     Acids  of  Group  III 148 

XVI.     The  Systematic  Procedure  of  Analysis  .     .  151 

NOTES        163 

INDEX        175 

ix 


CHEMICAL   ANALYSIS 

Part  I  —  Analytical  Operations 


CHAPTER   I 
INTRODUCTION 

The  science  of  chemistry  is  commonly  subdivided,  for 
purposes  of  convenience  in  reference  and  teaching,  into 
several  tolerably  distinct  branches.  The  usual  classifi- 
cation is  into  the  main  divisions  of  inorganic  and  organic 
chemistry,  each  of  which  may  in  turn  be  further  divided 
into  descriptive^  theoretical,  and  analytical  chemistry. 
Furthermore,  analytical  chemistry  may  itself  be  sepa- 
rated into  the  subdivisions  of  qualitative  and  quantitative 
analysis  ;  the  former  having  for  its  object  the  detec- 
tion of  chemical  elements  and  compounds,  and  the 
latter  the  relative  proportions  of  such  substances. 
Analytical  chemistry  is  commonly  taught  as  a  dis- 
tinct branch,  but  it  is  not  independent  of  the  other 
divisions  of  the  science  ;  and  hence,  in  all  discussions 
in  this  book,  both  as  to  theory  and  manipulation,  the 
presumption  is  that  the  student  has,  in  the  beginning 
of  the  course,  a  fair  knowledge  of  the  elements  of 
general  chemistry. 

It  obviously  is  essential  to  success  in  analysis  that 
the  analyst  should  have  a  clear  idea  of  the  operations 

1 


22  CHEMICAL  ANALYSIS 

involved  in  his  work,  as  well  as  of  the  compounds 
with  which  he  is  dealing  ;  and  therefore,  though  both 
manipulation  and  theory  are  assumed  to  have  been 
studied-  to  some  extent  in  connection  with  general 
chemistry,  it  is  deemed  well  to  review  many  of  the 
ordinary  operations  from  the  analytical  standpoint. 
The  first  part  of  this  book  is  devoted  largely  to 
such  a  review  ;  and  it  is  earnestly  recommended  that 
it  be  studied  closely,  and  that  all  of  the  experiments 
there  given  be  carefully  performed.  It  is  true  that 
the  time  spent  on  this  preliminary  work  will  delay 
somewhat  the  beginning  of  actual  analysis ;  but  it  is 
believed  that  the  student  will  be  repaid  in  the  end  by 
the  acquisition  of  a  clearer  conception  of  the  work  and 
of  more  skill  in  the  manipulation  of  apparatus. 

It  should  be  remembered  that  it  is  far  easier  to  form 
good  habits  than  to  correct  bad  ones ;  and  so  from  the 
beginning  the  attention  of  the  student  should  contin- 
ually be  directed  to  the  importance  of  the  following 
details  which,  though  simple  and  apparently  insignifi- 
cant, are  absolutely  essential  to  continued  success  in 
analysis. 

CARE   OF  APPARATUS 

(a)  Keep  all  apparatus  clean.  This  can  best  be  done 
by  cleaning  the  desk  and  apparatus  before  leaving  for 
the  day.  Of  course  this  does  not  apply  to  apparatus 
connected  with  unfinished  experiments. 

(b)  When  vessels  containing  materials  of  unfinished 
experiments  are  to  be  set  aside,  they  should  be  properly 
labeled. 


INTRODUCTION  3 

(c)  Provide  towels,  clean  rags,  soap,  and  a  covering 
for  the  clothes,  either  a  long  apron  or  a  workingman's 
overalls. 

(d)  Have  a  place  for  all  reagents  and  apparatus,  and 
keep  them  in  their  place.  Reagents  for  general  use 
should  not  be  kept  at  the  individual  desks.  This  is 
a  source  of  great  annoyance  and  injustice  to  one's 
neighbors. 

(e)  Use  all  care  in  keeping  the  reagents  pure. 
Stoppers  should  not  be  placed  on  the  desk  while  using 
the  bottles,  but  held  between  the  fingers.  No  foreign 
objects  should  be  dipped  into  the  bottles,  nor  should 
any  excess  of  reagents  be  poured  back  into  the 
bottles. 

(/)  Use  small  quantities  of  reagents.  It  is  best  to 
add  liquid  reagents,  drop  by  drop,  with  frequent 
shaking  of  the  test-tube,  so  that  secondary  reactions 
can  be  observed. 

LABORATORY  NOTES 

Provide  a  well-bound  notebook  for  the  subject  and 
use  it  for  nothing  else.  Keep  accurate  and  methodical 
records  of  all  experiments  performed.  These  records 
should  be  made  during  or  immediately  following  the 
performance  of  the  experiment,  and  not  transferred  or 
erased  afterwards.  Original  notes  of  an  unsuccessful 
experiment  are  more  valuable  than  a  well-written 
description  of  a  successful  experiment,  if  the  latter 
is  composed  in  the  absence  of  the  experiment. 

Some  states  prescribe  by  law  that  chemists,  in  giving 
expert  testimony  before  the  courts,  shall  present  only 


4  CHEMICAL  ANALYSIS 

such   data   as    are   recorded   in   the   presence    of    the 
experiments. 

If  desirable  the  original  notes  may  be  written  on 
alternate  lines  or  pages,  and  other  notes  of  interpre- 
tation added  at  leisure.  But  the  latter  should  be 
recorded  with  differently  colored  ink,  or  otherwise 
distinguished,  in  order  that  the  original  notes  be  not 
confused  with  subsequent  additions. 


CHAPTER   II 

THEORY  OF  ANALYTICAL  OPERATIONS 

Nature  of  Anal3rtical  Chemistry.  —  Analytical  chemistry 
has  already  been  defined  as  the  art  of  recognizing  the 
elements,  or  compounds,  which  may  be  present  in  any 
substance  ;  and,  as  the  nature  of  the  art  implies,  it 
commonly  is  practiced  upon  mixtures  of  one  kind  or 
another.  Such  mixtures  may  be  mechanical  only;  and 
in  such  a  case,  if  the  elements  of  the  mixture  are 
sufficiently  characterized  by  their  color,  crystalline 
form,  or  other  external  properties,  it  may  be  possible 
to  identify  or  even  to  separate  them  by  purely  mechan- 
ical means.  But  the  mixtures  with  which  the  chemist 
most  commonly  has  to  deal  are  those  in  which  the 
strictly  mechanical  element  plays  a  minor  part.  Such 
mixtures  are  produced  when,  by  any  appropriate  means, 
two  or  more  substances  are  brought  into  such  intimate 
contact  that  they  interpenetrate  each  other  even  to 
their  minutest  particles  —  the  molecules.  We  have 
examples  of  mixtures  of  this  class  in  the  air,  which 
practically  is  a  homogeneous  mixture  of  its  constituent 
gases  and  vapors  ;  in  common  "  solutions^"  such  as 
are  produced  when  any  suitable  material,  like  salt  or 
sugar,  is  treated  with  some  liquid  which,  like  water, 
has  the  power  of  "dissolving"  the  material  in.  ques- 
tion ;  or  in  alloys,  which  are  produced  when  two  or 

5 


6  CHEMICAL  ANALYSIS 

more  metals  are  united  by  fusion  into  a  mass  which  is, 
at  all  points,  of  uniform  composition.  Of  these  mix- 
tures, the  commonest  are  the  solutions  ;  and  these  are 
so  important,  from  the  standpoint  of  the  analytical 
chemist,  that  it  is  desirable  to  spend  some  time  in  a 
careful  study  of  their  properties. 

Solution.  —  In  a  general  sense  a  solution  is  the  prod- 
uct of  the  homogeneous  absorption  of  a  gas  by  a  gas, 
or  of  a  gas  by  a  liquid,  or  of  a  liquid  by  a  liquid,  or 
of  a  solid  by  a  liquid;  and  in  recent  years  the  term 
"  solid  solution "  has  been  applied  to  certain  homo- 
geneous solid  mixtures  of  which  the  alloy  mentioned 
above  may  serve  as  the  type.  But  specifically,  in 
speaking  of  a  solution,  we  have  in  mind  the  liquid 
product  of  the  absorption  by  a  liquid,  called  the 
solvent,  of  a  gas,  a  liquid,  or  a  solid,  called  the 
solute. 

It  has  been  found  of  all  gases,  and  of  some  liquids, 
that  they  are  capable  of  mixing  homogeneously  with 
one  another  in  all  proportions  ;  but,  on  the  contrary,  it 
has  not  been  found  possible,  under  ordinary  conditions, 
to  dissolve  a  gas  or  a  solid  in  a  liquid  in  any  desired 
proportion.  Sooner  or  later  a  point  is  reached  where 
the  solvent  refuses  to  take  up  more  of  the  solute ;  and 
at  this  point  the  solution  is  said  to  be  saturated.  In 
most  cases  the  application  of  heat  to  a  saturated  solu- 
tion will  enable  it  to  absorb  more  of  the  solute;  and 
the  application  of  cold  will  usually  result  in  the  sepa- 
ration of  a  part  of  the  material  already  dissolved.  In 
such  cases  we  may  recover, a  portion  of  the  solute  by 
the  mere  chilling  of  its  saturated  solution;  and  in  cases 


THEORY  OF  ANALYTICAL    OPERATIONS  7 

where  the  solute  is  practically  as  soluble  at  low  tem- 
peratures as  at  high  ones,  we  may  reach  the  same  end 
by  removing  a  part  or  the  whole  of  the  solvent  by 
evaporation.  It  may  be  mentioned  at  this  point  that 
we  have  still  another  means  of  separating  the  solute 
from  its  solution;  viz.^  by  the  addition  to  the  solu- 
tion of  some  material  which  will  decrease  the  solubility 
therein  of  the  solute,  without  changing  the  identity  of 
the  latter.  This  process  of  separation  is  of  considerable 
practical  importance,  and  we  shall  presently  have  occa- 
sion to  refer  to  it  again. 

Experiment  1 

(a)  Dissolve  5  grams  of  potassium  nitrate  in  25  c.c.  of  dis- 
tilled water,  at  a  temperature  of  15°-25°C.  Then  add  succes- 
sive portions  of  1  gram  each,  shaking  after  each  addition  until 
all  has  dissolved  that  the  solution  will  hold  at  this  temperature. 
Note  the  total  amount  added  and  then  raise  the  temperature  of 
the  solution  to  about  G0°,  —  as  hot  as  the  hand  can  bear  without 
too  much  discomfort,  —  and  add  more  of  the  finely  powdered 
salt  while  keeping  the  solution  from  cooling.  Note  the  extra 
amount  which  is  needed  at  this  temperature  to  saturate  the 
solution.  Now  cool  the  solution  quickly  and  note  the  result. 
Compare  any  material  which  may  separate  with  potassium  nitrate. 

(b)  Dissolve  5  grams  of  common  salt  in  25  c.c.  of  distilled 
water  at  15°-25°.  Now  add  successive  portions  of  ^  gram, 
shaking  after  each  addition  until  the  solution  is  saturated. 
Note  the  total  amount  dissolved.  Raise  the  temperature  as  in 
the  preceding  part  of  the  experiment,  and  see  whether  it  is  pos- 
sible to  dissolve  more  salt  in  the  hot  solution.  Allow  any  un- 
dissolved material  to  settle,  and  then  pour  off  some  of  the  clear 
solution  into  a  clean  dry  test-tube,  and  cool  as  much  as  pos- 
sible. Note  the  result.  Evaporate  a  portion  of  this  solution 
and  compare  the  residue  with  salt. 


8  CHEMICAL  ANALYSIS 

(c)  To  about  25  c.c.  of  a  clear  saturated  solution  of  common 
salt  add  50  c.c.  of  concentrated  hydrochloric  acid,  stirring  all 
the  time.  Note  the  result,  allowing  the  mixture  to  stand  for 
some  minutes.  Pour  off  the  clear  liquid  from  any  material 
which  may  have  separated,  and  press  a  little  of  the  latter  be- 
tween filter  papers,  to  remove  the  acid  liquor.  It  will  be  well 
to  remove  the  last  traces  of  acid  by  washing  the  residue  with  a 
little  saturated  brine.     Compare  the  residue  with  common  salt. 

It  will  have  been  seen,  in  the  performance  of  these 
experiments,  that  the  recovered  solute  is  of  the  same 
character  as  the  original  solute.  But  there  are  forms 
of  solution  in  which  this  is  not  the  case. 

Experiment  2 

Dissolve  a  small  piece  of  zinc  in  dilute  hydrochloric  acid  and 
evaporate  the  solution  to  dryness.  Compare  the  residue  with 
metallic  zinc. 

Solution  of  this  kind  may  be  called  chemical  solution,  in  dis- 
tinction from  the  simple  solutions  of  Exp.  1.  It  will  be  seen 
that  it  involves 

(1)  a  compound  solvent  —  HCl  +  water  —  which  itself  is  a 
simple  solution  ;  and  a  solute,  Zn  ; 

(2)  a  chemical  reaction  between  the  solute,  Zn,  and  one  con- 
stituent of  the  solvent,  HCl  —  Zn  +  2  HCl  =  ZnClg  +  2  H  —  in 
which  reaction  the  identities  of  the  solvent  and  of  the  solute 
are  changed ; 

(3)  a  simple  solution,  —  ZnCU  +  water. 

Chemical  solution  is  usually  the  result  of  the  mutual  reaction 
between 

(1)  an  acid,  or  a  base,  and  a  metal ; 

(2)  an  acid,  or  a  base,  and  a  salt; 

(3)  an  acid  and  a  base. 

But  it  may  happen,  as  when  metallic  sodium  is  dissolved  in 
water,  that  the  phenomenon  cannot  be  classified  under  any  of 
these  heads. 


THEORY  OF  ANALYTICAL    OPERATIONS  9 

Simple  solution  is  often  a  necessary  predecessor  of 
chemical  solution,  as  has  been  seen  in  Exp.  2 ;  and,  in 
general,  it  prepares  the  way  for  chemical  action  by 
placing  the  reagents  in  close  contact. 

Experiment  3 

Mix  .5  gram  of  dry  potassium  iodide  with  .5  gram  of  dry 
mercuric  chloride  in  a  dry  mortar,  and  rub  the  mixed  salts  well 
together  with  the  pestle.  Note  the  result.  Add  a  little  water 
and  rub  again. 

Furthermore,  simple  solution  may  be  necessary  to  the 
continuance  of  chemical  action,  in  order  that  the  products 
of  reaction  may  be  removed  from  between  the  reagents. 

Experiment  4 

Add  a  bit  of  zinc  to  5  c.c.  of  concentrated  sulphuric  acid  in  a 
test-tube ;  leave  for  a  few  moments,  noting  all  that  happens. 
Now  transfer  the  contents  of  the  tube  to  a  dish  containing  15- 
20  c.c.  of  water. 

When  zinc  is  treated  with  conceutrated  sulphuric  acid,^  chemi- 
cal action  occurs  for  a  short  time  only,  and  then  ceases  entirely. 
The  explanation  is  probably  this :  zinc  sulphate,  insoluble  in 
concentrated  sulphuric  acid,  coats  the  zinc  and  prevents  further 
contact  of  the  reagents.  The  addition  of  water,  in  which  zinc 
sulphate  is  very  soluble,  removes  the  coating  and  permits  chemi- 
cal action  to  go  on  once  more. 

Properties  of  the  Solute. — So  far  in  our  study  of  the 
phenomena  of  solution,  we  have  considered  only  those 
properties  of  the  solute  which  are  associated  with  its 
solid  condition,  —  when,  in  point  of  fact,  it  cannot 
properly  be  called  a  solute.  Let  us  now  see  whether 
we  can  discover  anything  concerning  the  properties 
of  the  true  solute,  —  the  body  in  solution. 


10  CHEMICAL  ANALYSIS 

We  have  seen  that  a  solution  which  is  saturated 
with  a  given  body  at  one  temperature  may  acquire  the 
power  of  dissolving  an  additional  quantity  of  that  body 
in  consequence  of  an  elevation  of  temperature,  and 
that,  on  the  contrary,  it  may  give  up  a  portion  of  its 
solute  if  its  temperature  is  lowered.  That  is  to  say, 
if  we  have  a  "  system  "  consisting  of  a  limited  quantity 
of  some  saturated  solution  in  contact  with  an  excess  of 
its  solute,  there  will  be  for  any  given  temperature  a 
concentration  of  the  solution  at  which  there  will  be 
a  condition  of  equilibrium  between  the  dissolved  and 
undissolved  solute.  This  condition  is  entirely  analo- 
gous to  that  which  is  observed  when  a  volatile  liquid 
is  exposed  in  contact  with  a  limited  volume  of  air  or 
other  gas.  In  the  latter  case  the  liquid  will  volatilize, 
—  rapidly  at  first,  and  afterwards  more  slowly,  —  until 
the  concentration  of  its  vapor  in  the  atmosphere  to 
which  it  is  exposed  has  reached  a  certain  limit  which 
will  be  dependent  on  the  temperature.  With  a  rise  in 
temperature,  more  liquid  will  pass  into  the  state  of 
vapor;  with  a  fall,  a  portion  of  the  liquid  already 
vaporized  will  be  condensed  again. 

This  analogy  has  been  recognized  for  many  years  ; 
but  it  is  now  hardly  more  than  a  decade  since  first  its 
completeness  was  fully  demonstrated. 

Colloids  and  Crystalloids.^  —  It  had  been  shown  by 
Graham  (1842)^  that  certain  colloid  solutes,  whose  solu- 
tions are  not  real  liquids,  but  emulsions,  cannot  pass 
through  porous  membranes  ^  —  such  as  parchment  — 
and  that  most  crystalloid  solutes,  whose  solutions  are 
real  liquids,   readily  penetrate  such  septa.      He  first 


THEORY   OF  ANALYTICAL    OPERATIONS  11 

put  separate  solutions  of  a  colloid  and  a  crystalloid 
into  separate  open  cylinders  whose  bottoms  were  closed 
with  parchment,  and  then  suspended  the  cylinders  in 
vessels  of  water  so  that  the  membranes  were  immersed. 
After  a  few  hours  a  large  part  of  the  crystalloid  had 
passed  through  the  parchment  into  the  water  in  the 
outer  vessel;  and  by  renewing  this  water  all  of  the 
crystalloid  was  finally  extracted  from  the  cylinder. 
From  the  other  cylinder,  however,  no  colloid  had 
passed  out. 

Osmosis. — Pfeffer,^  the  botanist,  in  demonstrating  and 
measuring  the  internal  bursting  force  of  plant  cells 
(1877),  established  the  fact  that  crystalloids,  though  they 
do  not  pass  through  the  so-called  "  semi-permeable " 
membranes, — of  which  protoplasm  ^  is  a  type, — do  press 
strongly  against  the  partition  in  their  futile  attempt  to 
penetrate  it.  Connecting  a  mercury  gauge  and  ther- 
mometer with  a  membrane,  composed  of  a  porous  cell 
coated  with  copper  ferrocyanide,  and  charging  this  ap- 
paratus with  saccharine  solutions  of  different  strengths, 
he  found  that  different  concentrations  of  solution  pro- 
duced correspondingly  different  pressures  within  the 
apparatus  when  the  temperature  was  kept  constant, 
and  that  for  any  given  concentration  the  pressure  varied 
as  the  absolute  temperature.  He  showed,  therefore, 
that  the  relations  of  concentration,  pressure,  and  tem- 
perature, which  are  shown  by  sugar  in  its  solutions,  are 
identical  with  those  manifested  by  gases, — of  which  it 
will  be  remembered  that  the  concentration  or  density 
of  a  given  mass  varies  directly  as  the  pressure  and 
inversely  as   the   absolute  temperature.     To  the  form 


12  CHEMICAL   ANALYSIS 

of  tension  exercised  by  the  dissolved  sugar  he  gave  the 
name  osmotic  pressure. 

Law  of  Osmotic  Pressure. — Van't  Hoff^  (1887)  found 
that  a  large  number  of  solutions  behave  like  that  of 
sugar,  and  announced  the  following  law:  The  osmotic 
pressure  of  a  substance  in  solution  is  identical  with  the 
pressure  which  it  would  exert  were  it  in  the  form  of  a 
gas  occupying  the  same  volume  {i.e.^  the  volume  of  the 
solution)  at  the  same  temperature.'^ 

We  may  conveniently  express  the  simple  law  which 
governs  the  phenomena  of  gas  and  osmotic  pressures  in 
the  following  form :  — 

^,     MT  MT 

y= or  !)  =  ——-» 

p  ^        V 

wherein  M  represents  the  number  of  molecules  ^  in  a 
given  body  of  gas,  T  and  p  the  temperature  and  pres- 
sure, and  V  the  volume.  Certain  gases,  such  as  oxygen, 
nitrogen,  and  hydrogen,  are  obedient  to  this  law  within 
very  wide  limits;  but  there  are  vapors  whose  behavior 
with  regard  to  it  is  apparently  anomalous.  Evidently 
V  can  be  made  constant,  and  T  and  p  can  be  measured 
with  any  desired  degree  of  accuracy.  And  therefore 
unless  there  can  be  a  change  in  the  value  of  M^  any 
change  in  T  ought  to  be  accompanied  by  an  exactly 
proportional  change  in  p.  Now  we  find  that  certain 
vapors  —  such  as  that  of  ammonium  chloride  —  give 
greater  pressures  than  can  be  accounted  for  by  either 
the  value  of  T,  or  the  value  of  il[f  which  is  based  upon  the 
commonly  accepted  molecular  weight;  and,  as  has  been 
indicated,  we  find  the  explanation  of  this  behavior  in 


THEORY   OF  ANALYTICAL    OPERATIONS  13 

the  fact  that  the  molecule  NH^Cl  is  split  up,  or  "  disso- 
ciated," 1  when  we  seek  to  vaporize  it,  into  the  smaller 
molecules  NHg  and  HCl.  The  analogies  between  the 
behavior  of  gases  and  substances  in  solution  seem  to 
extend  to  this  phenomenon  of  dissociation,  for  it  has 
been  observed  of  many  solutes  that  their  osmotic  pres- 
sures are  so  large  as  to  be  accounted  for  only  on  the 
supposition  that  their  molecules  are  split  up  in  solution 
and  thereby  increased  in  number.  Sugar  and  other 
bodies  of  its  neutral  character  obey  the  simple  law  as 
stated  above;  but  acids,  bases,  and  salts  in  aqueous 
solution  usually  exhibit  anomalous  pressures. 

Freezing  Point  Depression.  —  Moreover,  this  is  not  the 
only  evidence  which  bears  upon  the  question  of  the 
dissociation  of  the  molecules  of  solutes.  It  is  a  matter 
of  common  knowledge  that  the  boiling  and  freezing 
points  of  aqueous  solutions  are  respectively  higher  and 
lower  than  those  of  pure  water.  These  relations  were 
studied  carefully  by  Raoult,^  who  showed  that  the  phe- 
nomenon is  a  general  one  and  that :  — 

(a)  When  any  substances  are  dissolved  in  inactive 
solvents,  the  changes  in  the  freezing  and  boiling  points 
of  the  solvents  vary  with  the  amounts  of  substance 
dissolved. 

(b)  When  equal  weights  of  different  substances  are 
dissolved  in  equal  amounts  of  the  same  solvent,  the 
changes  vary  inversely  with  the  molecular  weights  of 
the  solutes. 

It  was  found  of  many  bodies  —  such  as  sugar  —  that 
equal  depressions  of  the  freezing  point  were  pro- 
duced by  the  solution  of  equimolecular  proportions  in 


14  CHEMICAL  ANALYSIS 

water;  and  in  such  cases  the  depressions  were  exactly 
in  inverse  ratio  to  the  molecular  weights.  In  other 
cases,  however,  the  solutions  of  equimolecular  weights 
of  different  substances  produced  unequal  depressions; 
and  the  solution  of  different  weights  of  a  given  sub- 
stance produced  depressions  which  were  not  in  exact 
ratio  to  the  weights  so  dissolved.  In  the  latter  anoma- 
lous cases  the  depressions  were  greater  than  seemed  to 
be  called  for  by  the  amount  of  matter  which  had  been 
dissolved,  as  naturally  would  be  the  case  if  the  mole- 
cules of  the  dissolved  substances  were  dissociated  into 
more  numerous  and  smaller  molecules;  and  the  sub- 
stances which  exhibited  this  behavior  were  those  which 
show  abnormal  osmotic  pressures,  namely,  the  majority 
of  acids,  bases,  and  salts. 

In  these  two  pieces  of  independent  evidence  we 
have  a  strong  demonstration  of  the  fact  that  many  sub- 
stances exhibit,  when  dissolved  in  water,  a  peculiar 
structural  condition  in  which  their  molecules  are  split 
up  into  smaller  bodies  than  are  indicated  by  their 
accepted  formulae;  and  we  have  to  inquire  what 
further  evidence  we  have  which  will  throw  light  upon 
the  precise  nature  of  these  submolecules.  We  shall 
find  this  evidence  in  connection  with  the  behavior  of 
solutions  which  are  subjected  to  the  passage  of  an 
electric  current. 

Electrolytes.  —  It  lias  long  been  known  that  the  con- 
ductivity exhibited  by  liquids  is  unlike  that  of  metallic 
conductors,  in  that  the  latter  are  not  affected  chemi- 
cally by  the  passage  of  a  current,  whereas  the  former 
are  decomposed  with  separation  at  the  electrodes  — 


THEORY  OF  ANALYTICAL    OPERATIONS  15 

the  points  where  the  current  enters  and  leaves  the 
liquid- — of  products  of  varying  character.  In  1834 
Faraday  ^  suggested,  in  explanation  of  this  phenomenon, 
that  the  liquid  which  conducts  electricity  has  in  solu- 
tion a  compound  whose  molecules  are  divided  into 
freely  moving  particles,  some  of  which  are  charged 
with  positive  and  the  rest  with  negative  electricity. 
He  named  such  compounds  electrolytes;  and  to  the 
hypothetical  fragments  of  their  molecules  he  gave  the 
name  of  ions.  Those  which  were  assumed  to  be  posi- 
tively charged  were  called  cations,  and  were  either 
metals,  or  atom-complexes,  like  NH^,  which  react  analo- 
gously to  metals.  Those  bearing  a  negative  charge 
were  termed  anions,  Sind  were  such  bodies  as  the  halo- 
gens and  acid  radicles.  The  attraction  or  neutralizing 
effect  which  ions  of  opposite  polarities  were  supposed 
to  exercise  upon  each  other,  was  held  to  maintain  the 
identity  of  the  solute  until  the  solution  was  subjected 
to  the  passage  of  an  electric  current;  whereupon  the 
introduction  of  electrodes  of  opposite  polarities  upset 
the  equilibrium  previously  existing  between  the  ions 
and  caused  them  to  migrate,  —  the  negative  ions  going 
toward  the  positive  electrode,  and  the  positive  ions  in 
the  opposite  direction.  The  appearance  of  decomposi- 
tion products  at  the  electrodes  was  explained  as  being 
due  to  the  union  of  the  ions,  upon  arrival  at  those 
points,  to  the  molecular  condition  or  to  compounds 
with  the  elements  of  water. 

In  1887  it  was  demonstrated  by  Arrhenius^  that  the 
solutions  which  exhibit  normal  osmotic  pressures  and 
freezing  point  depressions  are  nonconductors  of  elec- 


16  CHEMICAL  ANALYSIS 

tricity,  and  that  their  solutes  are  not  electrolytes. 
Conversely,  the  solutions  which  give  abnormal  osmotic 
pressures  were  proved  to  contain  ionized  solutes ; 
and  it  was  shown,  furthermore,  by  highly  accurate 
experimental  methods,  that  the  degree  of  their  con- 
ductivity is  proportional  to  the  amount  of  dissocia- 
tion as  measured  by  the  osmotic  pressure.  Between 
the  extremes  presented  by  bodies  like  sugar,  which 
are  characterized  by  little  chemical  reactivity  and  the 
absence  of  conductivity  and  dissociation,  and  such 
substances  as  salts  and  strong  acids  and  bases,  which 
are  distinguished  by  great  reactivity  and  perfect  con- 
ductivity and  dissociation,  were  arranged  the  other 
varieties  of  chemical  compounds,  which  possess  various 
but  proportional  activities  of  the  three  kinds. 

With  the  establishment  of  these  facts  the  phenom- 
enon of  electrolytic  dissociation  received  a  new  signifi- 
cance from  the  standpoint  of  analytical  chemistry.  The 
behavior  of  molecules  in  solution  was  seen  to  be  chiefly 
dependent  upon  their  tendency  toward  or  from  disso- 
ciation. The  solutions  of  strongly  ionized  bodies  are 
characterized  rather  by  the  reactions  of  the  ions  than 
by  the  properties  of  the  undissociated  molecules.  In 
the  case  of  sodium  chloride,  for  example,  the  solution 
presents  certain  definite  properties  which  are  charac- 
teristic of  the  chlorine  and  sodium  ions,  and  practically 
none  which  are  characteristic  of  salt  itself.  In  the  case 
of  sugar  solutions,  on  the  contrary,  such  properties  as 
are  manifested  are  those  of  the  sugar  molecule  alone ; 
and  no  indication  is  to  be  seen  in  them  of  the  nature 
of  the  constituent  elements  of  sugar. 


THEORY  OF  ANALYTICAL    OPERATIONS  17 

Analytical  Significance  of  Ions.  —  Borrowing  an  illustra- 
tion from  Ostwald,  let  us  assume  that  we  have  to  deal 
with  50  basic  and  50  acidic  units  of  some  kind,  which 
may  in  theory  unite  to  form  2500  distinct  compounds 
with  as  many  sets  of  distinctly  individual  properties. 
Were  the  analyst  compelled  to  recognize  these  com- 
pounds singly,  in  the  solid  condition,  he  obviously 
would  have  to  be  familiar  with  the  properties  of  each 
individual  among  the  whole  number;  and  were  he  to 
attempt  to  identify  the  individuals  that  might  be 
present  in  a  mixture,  the  task  would  be  beyond 
accomplishment.  Were  the  compounds  not  dissoci- 
able in  solution,  his  problem  would  still  be  scarcely 
less  difficult  of  solution  ;  but,  being  dissociable,  his 
task  is  made  comparatively  light.  Since  the  proper- 
ties of  the  solution  of  an  ionized  compound  are  merely 
the  sum  of  the  properties  of  its  ions,  and  since  the 
total  number  of  ions  with  which  we  have  assumed  it 
necessary  to  deal  is  100,  it  follows  that  the  knowledge 
of  100  sets  of  properties  is  sufficient  for  the  identifica- 
tion of  any  of  the  2500  compounds.  If,  as  it  some- 
times happens,  the  substance  under  examination  is  not 
soluble  or  readily  dissociated,  the  analyst  has  only  to 
convert  it  by  appropriate  means  into  a  body  which 
18  soluble  and  dissociable,  and  then  to  determine  its 
nature  from  the  character  of  the  latter  substance. 

Laws  of  Electrolytic  Dissociation.  —  So  far  we  have  con- 
sidered only  the  qualitative  effects  of  electrolytic  dis- 
sociation; let  us  now  examine  briefly  the  quantitative 
effects,  which  are  of  no  less  importance  to  the  analytical 
chemist. 


18  CHEMICAL  ANALYSIS 

As  has  been  said  already,  different  electrolytes  have 
been  found  to  show  great  dissimilarity  in  conductivity 
and  ionization,  even  when  dissolved  in  equimolecular 
proporjtions.  But  it  also  has  been  found  that  all  are 
obedient  to  the  same  law  with  regard  to  the  degrees 
of  their  dissociation,  and  that  the  dissimilarities  are 
accounted  for  by  constants  which  depend  upon  the 
nature  of  the  electrolytes.  The  observed  relations 
between  the  amounts  of  dissociated  and  undissociated 
electrolyte  in  a  solution  are  expressed  most  simply  for 
binary  compounds  in  the  equation 

a.b  =  k.c, 

wherein  a  represents  the  concentration  of  the  positive 
ions,  b  that  of  the  negative  ions,  e  that  of  the  mole- 
cules of  undissociated  material,  and  k  a  constant  func- 
tion of  the  electrolyte.  By  assuming  a  value,  such  as 
unity,  for  the  total  amount  of  electrolyte  in  solution, 
and  by  representing  the  amount  of  dissociated  material 
by  «;,  and  the  volume  of  the  solution  by  v,  we  may 
expand  this  equation  to  a  somewhat  more  instructive 
form  :  — 

c,  concentration  of  undissociated  electrolyte  = ; 

V 

a  and  5,  concentrations  of  the  two  ions, —  either  ion  =  -• 

V 

By  substitution  we  obtain  the  equation  in  the  form 

=  kv. 


1-a 

Inspection  of  these  equations,  which  are  merely  the 
formal  expression  of  observed  fact,  reveals  :  — 


THEORY  OF  ANALYTICAL    OPERATIONS  19 

(1)  that  increase  in  a  (or  b)  will  be  accompanied  by 

a.h 
an  increase   in  the  ratio   — ?    i.e..   the  free  ions   will 

0 

increase  and  the  molecules  will  decrease ; 

(2)  that  decrease  in  a  (or  h)  will  have  the  opposite 
effect,  i.e.,  the  free  ions  will  decrease  and  the  mole- 
cules will  increase; 

(3)  that  the  degree  of  dissociation  may  vary  in  either 
direction  according  as  k  is  increased  or  decreased  by 
variation  in  the  nature  of  the  electrolyte ; 

(4)  that  dilution  of  a  solution,  and  corresponding 
increase  of  v,  will  call  for  an  increase  in  the  propor- 
tion of  dissociated  solute,  the  degree  of  dissociation 
approaching  totality  as  its  limit,  as  the  dilution  is 
indefinitely  increased; 

(5)  that  concentration  will  have  the  opposite  effect, 
and  that  the  ratio  of  dissociated  to  undissociated  solute 
will  reach  its  minimum  limit  in  a  saturated  solution. 

Further  inspection  of  the  equation  a.b  =  k.c  will  reveal 
another  fact  which  is  of  great  practical  significance 
for  the  analytical  chemist.  It  is  evident  that  in  the 
solution  of  any  given  electrolyte,  at  a  fixed  tempera- 
ture, the  only  possible  variants  will  be  a,  5,  and  c. 
Let  us  suppose  that  it  is  possible  in  some  way  to  intro- 
duce an  added  quantity  of  one  ion,  so  that  either  con- 
centration a  ov  h  will  be  increased.  This  being  done, 
the  increase  in  the  product  a.h  will  demand  an  increase 
in  the  value  c.  But  the  only  way  in  which  c  may  be 
increased  is  through  the  return  from  dissociation  of  a 
certain  proportion  of  the  ions.  Assuming  the  concen- 
tration h  to  have  been  increased,  the  concentration  a 


20  CHEMICAL  ANALYSIS 

must  be  diminished  until,  by  the  decrease  in  a.h  and 
the  corresponding  increase  in  c,  the  original  condition 
of  equilibrium  has  been  restored.  In  case  that  we  are 
dealing  with  a  saturated  solution  of  the  electrolyte, 
any  increase  in  c  will  result  in  supersaturation  of  the 
solution ;  and  we  shall  see  that  a  portion  of  the  solute 
may  separate  in  solid  form.  In  fact,  we  have  already 
seen  this  in  a  practical  way  in  Exp.  1,  c. 

In  the  saturated  sodium  chloride  solution  of  that 
experiment,  a  considerable  portion  of  the  solute  was 
present  in  the  form  of  Na  and  CI  ions;  and  the  re- 
mainder was  present  in  the  molecular  condition  in 
quantity  sufficient  to  produce  saturation.  The  addi- 
tion of  concentrated  HCl,  whose  solution  is  very 
strongly  dissociated,  introduced  a  very  large  excess 
of  CI  ions  in  the  salt  solution;  and,  in  consequence, 
the  reunion  of  sodium  and  chlorine  ions  to  the  molec- 
ular state  was  set  up  and  continued  until  equilibrium 
had  been  restored.  But  as  the  solution  had  already 
been  saturated  with  the  molecules  of  salt,  these  re- 
formed molecules  were  forced  to  separate  in  the  solid 
form. 

If  we  dissolve  together  two  substances  which  are  dis- 
sociated more  equally,  such  as  KCl  and  NaCl,  we  find 
that  less  action  of  this  sort  takes  place ;  but  when,  of 
our  two  solutes  with  a  common  ion,  one  is  more  strongly 
dissociated  than  the  other,  the  weaker  is  forced  back 
to  the  molecular  and  inactive  condition. 

The  constant  k  has  a  very  uniform  value  for  neutral 
salts,  but  varies  considerably  for  acids  and  bases,  being 
high  for  strong  acids  and  low  for  weak  ones. 


THEORY  OF  ANALYTICAL   OPERATIONS  21 

Regarding  the  dissociation  values  of  k,  Ostwald  has 
separated  acids,  bases,  and  salts  into  three  classes : 

Class  1 :  Neutral  salts,  strong  acids,  and  strong  bases. 
The  strong  acids  mentioned  are  hydrochloric,  hydro- 
bromic,  hydriodic,  nitric,  chloric,  and  sulphuric  ;  the 
strong  bases  are  hydroxides  of  the  alkali  and  alkali- 
earth  metals. 

Class  2 :  Moderately  strong  acids  and  bases.  The 
acids  are  phosphoric,  sulphurous,  and  acetic ;  the  bases 
are  the  hydroxides  of  ammonium,  silver,  and  magnesium. 

Class  3 :  Weak  acids  and  bases.  The  acids  are  car- 
bpnic,  hydrosulphuric,  hydrocyanic,  silicic,  and  boracic ; 
the  weak  bases  are  the  hydroxides  of  the  trivalent 
metals  and  of  those  divalent  metals  not  mentioned  in 
Classes  1  and  2. 

Applications.  —  This  discussion  of  the  theories  and 
laws  of  electrolytic  dissociation  enables  us  to  explain 
many  important  operations  and  reactions  in  analytical 
chemistry,  which  otherwise  could  hardly  be  understood. 

A  few  of  the  explanations  may  be  conveniently  formu- 
lated by  questions  and  answers : 

1.  How  does  ionization  aid  chemical  activity? 

By  dissociation  of  the  solute  into  its  ions,  making  it 
possible  for  them  to  combine  with  other  ions. 

2.  How  may  heat  aid  chemical  activity  ?  ^ 

By  producing  rapid  vibrations  of  the  molecules,  which 
thus  increases  the  speed  of  the  reaction. 

3.  How  may  dilution  aid  chemical  activity? 

By  expanding  the  volume,  thus  decreasing  the  pree- 
SLiro^  and  increasing  the  degree  of  dissociation. 


22  CHEMICAL  ANALYSIS 

4.  Why  is  the  activity  of  an  add  or  a  base  usually 
decreased  hy  adding  some  salt  of  that  acid  or  base  ? 

Two  examples  are  given  : 

(1)  ^The  addition  of  sodium  acetate  to  acetic  acid  decreases 
the  solvent  power  of  the  acid,  since  the  salt  is  more  strongly 
dissociated  than  the  acid,  and  causes  a  portion  of  the  latter  to 
reassume  the  molecular  condition  by  increasing  the  concentra- 
tion of  the  CgHgOg  ions. 

(2)  The  addition  of  ammonium  chloride  to  ammonia^  water 
decreases  the  solvent  action  of  the  latter  by  increasing  the  con- 
centration of  the  NH4  ions,  and  decreasing  the  dissociation  and 
activity  of  the  NH4OH. 

5.  When  an  excess  of  a  normal  salt  of  a  weak  acid  is 
added  to  a  solution  of  a  strong  acid,  why  is  the  activity 
of  the  strong  acid  destroyed,  and  that  of  the  resulting 
weak  acid  greatly  weakened? 

If  an  excess  of  sodium  acetate  is  added  to  a  solution 
of  calcium  phosphate  in  very  dilute  hydrochloric  acid, 
the  phosphate  will  be  precipitated  in  spite  of  the  fact 
that  it  is  soluble  in  both  hydrochloric  and  acetic  acids. 
The  explanation  of  this  behavior  is  as  follows :  ^  Hydro- 
chloric acid  and  sodium  acetate  react  to  form  sodium 
chloride  and  acetic  acid.  The  latter,  in  the  presence 
of  the  excess  of  sodium  acetate,  is  forced  back  into  the 
inactive  molecular  condition  in  which  it  is  no  longer 
able  to  hold  the  phosphate  in  solution. 

6.  Why  does  the  addition  of  a  solvent  having  an  ion 
in  common  with  that  of  a  solute  salt  tend  to  precipitate 
the  solute  ? 

This   question   already   has    been    answered   in   the 


THEORY   OF  ANALYTICAL    OPERATIONS  23 

explanation  of  the  precipitation  of  common  salt  from 
its  solution  by  the  addition  of  hydrochloric  acid. 

7.  Wh^  do  reagents  behave  differently  towards  the 
same  elements  in  different  compounds  f 

For  example,  hydrogen  sulphide  precipitates  black 
cupric  sulphide  from  a  solution  of  cupric  sulphate,  but 
not  from  a  solution  of  potassium  cuprous  cyanide. 
Another  example,  silver  nitrate  precipitates  white  silver 
chloride  from  a  solution  of  potassium  chloride,  but  not 
from  a  solution  of  potassium  chlorate.  The  general 
answer  to  the  question  is  that  the  chemical  activity  of 
a  compound  depends  on  its  dissociated  ions,  —  not  on 
the  presence  of  certain  elements.  Hydrogen  sulphide, 
HgS,  reacts  with  cupric  sulphate,  CuSO^,  because  the 
latter  is  ionized  into  Cu  and  SO4.  Hydrogen  sul- 
phide does  not  react  with  potassium  cuprous  cyanide, 
K3Cu(CN)4,  because  the  latter  gives  no  free  Cu  ions, 
but  the  molecule  is  dissociated  into  the  ions,  3K  and 
Cu(CN),. 

In  the  second  example  silver  nitrate,  AgNOg,  reacts 
with  potassium  chloride,  KCl,  because  the  latter  is 
dissociated  into  K  and  CI;  but  silver  nitrate  does  not 
react  with  potassium  chlorate,  KClOg,  as  the  latter  salt 
is  dissociated  into  K  and  CIO3. 

8.  Wh^  do  reagents  behave  alike  with  various  salts  of 
the  same  metal  f 

When  we  say  of  any  substance  that  it  is  a  salt  of 
a  certain  metal,  —  such  as  copper,  —  we  imply  that  it 
dissociates  in  solution  with  the  formation  of  ions  of 
that    metal.       These    always    react    alike,    no    matter 


24  CHEMICAL  ANALYSIS 

what   the   negative  ions   be  with  which  they  are   in 
equilibrium. 

9.  Whi/  are  normal  salts  usually  better  precipitants 
than  their  corresponding  acids  or  bases  f 

For  example,  calcium  chloride  readily  reacts  with 
ammonium  carbonate,  but  not  with  carbonic  acid.  The 
following  equations  illustrate  the  comparative  reactivi- 
ties of  normal  salts,  acid  salts,  and  acids : 

(NH4)2C03  +  CaCl2  yields  an  immediate  precipitate; 

H(NH4)C03  +  CaCl2  yields  a  tardy  precipitate; 

HgCOg  +  CaClg  yields  no  precipitate. 
Normal  salts  are  most  completely  dissociated,  while 
weak  acids  are  very  slightly  dissociated.  Acid  salts  of 
weak  acids  partake  of  the  nature  of  both  normal  salts 
and  weak  acids.  As  ammonium  carbonate  is  a  normal 
salt,  it  is  moBe  completely  dissociated  than  either  the 
acid  salt,  H(NH4)C03,  or  the  acid  HgCOg,  —  and  hence 
it  reacts  with  calcium  chloride  more  readily. 

10.  Why  does  an  excess  of  a  strong  basic  precipitant 
redissolve  many  precipitates  from  salts  of  weak  bases? 

For  example,  a  weak  solution  of  sodium  hydroxide 
precipitates  aluminum  hydroxide  from  a  strong  solution 
of  aluminum  sulphate,  but  on  adding  an  excess  of  the 
precipitant,  the  precipitate  disappears.  Two  reactions 
occur  here : 

(a)  Aluminum  hydroxide  is  formed : 

6NaOH  +  Al2(S04)3  =  2Al(OH)3  +  3Na2S04;  and  on 
adding  more  sodium  hydroxide  the  white  precipitate 
dissolves,  forming  sodium  aluminate,  — 

(b)  Al(OH)3  4-3NaOH  =  Na3A103  +  3H20. 


THEORY  OF  ANALYTICAL    OPERATIONS  25 

Interpreted  in  terms  of  the  ionic  theory,  aluminum 
being  a  very  weak  basic  metal,  its  hydroxide  is  easily 
influenced  by  a  strong  base.  In  aqueous  solution 
A1(0H)3  is  in  equilibrium,  being  partly  dissociated 
into  the  ions  A1+  and  3  0H~,  and,  by  loss  of  water, 
partly  into  the  ions  H"*"  and  A102~.  When  a  strong 
base  like  NaOH  is  added,  it  neutralizes  the  acid  HAlOg, 
forming  NagAlOg  and  water.  This  destroys  the  equi- 
librium, and  more  H+  and  A102~  are  developed,  only 
to  be  in  turn  neutralized  by  more  NaOH.  And  so 
the  process  continues  till  all  of  the  A1(0H)3  goes  into 
solution  as  NaoAlOo. 


CHAPTER  HI 

METHODS  OF  ANALYTICAL  SEPARATION 

Object  of  Separation.  —  It  is  only  in  rare  cases  that  the 
chemist  is  able  to  recognize  and  identify  individual 
elements  or  compounds  in  the  mixtures  which  contain 
them,  without  having  first  separated  them  from  the 
other  bodies  there  present.  In  some  cases,  —  the  mix- 
tures being  purely  mechanical,  —  a  mechanical  treatment 
is  sufficient  to  accomplish  the  separation;  in  other 
cases,  —  as  when  the  substances  are  present  in  solution, 
—  it  is  necessary  in  addition  to  make  use  of  chemical 
or  physical  processes,  by  which  means  the  material 
under  examination  is  converted  into  such  form  that 
the  recognition  of  its  elements  is  possible.  We  have 
therefore  two  classes  of  separations,  —  the  members  of 
the  first  class  being  of  a  mechanical  nature,  whereas 
those  of  the  second  are  of  either  physical  or  chemical 
character.  The  principal  separations  of  the  first  class 
are  brought  about  by  the  operations  of  deeantation^ 
filtration^  and  washing. 

MECHANICAL   SEPARATIONS 

Decantation.  —  When  we  have  a  mixture  of  a  solid 
with  a  liquid  in  which  it  is  insoluble,  or  a  mixture 
of  two  liquids  which  are  mutually  insoluble,  we  may 

26 


METHODS   OF  ANALYTICAL   SEPARATION         27 

separate  them  by  this  process,  provided  that  their  specific 
gravities  are  so  different  that  one  of  the  compounds  of  the 
mixture  will  settle  and  separate  completely  from  the  other. 
From  a  mixture  of  liquid  with  solid, — for  example,  water 
and  silver  chloride,  —  we  may  remove  most  of  the  liquid 
by  careful  pouring  or  by  suction  with  a  pipette.  From  a 
mixture  of  liquids,  —  such  as  water  and  ether,  —  we  may 
remove  either  layer  with  the  pipette,  or  we  may  draw  off 
the  lower  layer  by  means  of  a  separatory  funnel. 

Though  decantation  never  separates  completely,  it  is 
convenient  for  the  removal  of  the  bulk  of  liquids  from 
finely  divided  precipitates  which  pass  through  the  filter 
paper,  or  from  gelatinous  precipitates  which  clog  its 
pores.  Separation  can  be  hastened  by  centrifugal  shak- 
ing of  the  mixture  before  decantation. 

Ezperiment  5 

(a)  Dissolve  a  few  crystals  of  silver  nitrate  in  10  c.c.  of  water 
in  a  test-tube,  and  then  add  dilute  hydrochloric  acid,  drop  by 
drop,  until,  by  shaking,  the  white  silver  chloride  settles  beneath  a 
clear  liquid.  Decant  the  liquid  by  pouring  it  oif  with  a  glass 
rod  held  against  the  edge  of  the  test-tube.  Add  more  water 
to  the  solid  and  decant  again  by  immersing  the  tip  of  a  pipette 
in  the  clear  liquid  and  sucking  it  off  with  the  mouth  (never  allow 
the  liquid  to  rise  to  the  mouth).  Close  the  mouth-end  of  the 
pipette  with  the  tongue,  lift  out  the  pipette,  and  when  the  tongue 
is  removed  the  liquid  will  flow  out. 

(b)  Mix  5  c.c.  each  of  ether  and  water  in  a  test-tube  by  shak- 
ing vigorously.  The  lighter  ether  will  rise  to  the  top.  Remove 
either  the  ether  or  water  with  a  pipette. 

Filtration.  —  Filtration  is  the  separation  of  a  solid 
residue   from   a   liquid  filtrate   by   means   of  a  porous 


28  CHEMICAL  ANALYSIS 

partition  impervious  to  the  residue.  The  partition 
most  frequently  used  in  analytical  work  is  unsized 
paper,  supported  in  a  glass  funnel.  A  circular  paper 
is  folded  twice^  so  as  to  form  the  quadrant  of  a  circle, 
and  is  then  fitted  into  a  glass  funnel  and  dampened,  so 
as  to  adhere  closely  to  the  sides  of  the  funnel.  For 
rapid  filtration  it  is  convenient  first  to  fold  the  paper 
once  across  the  middle,  and  then  to  "plait"  it  on 
radial  lines,  so  that  it  resembles  finally  a  folding  paper 
fan.  On  opening  a  paper  so  folded,  it  will  be  seen  to  fit 
loosely  in  the  funnel  and  to  leave  numerous  channels 
through  which  the  filtrate  may  escape. 

To  prevent  spattering,  the  beveled  tip  of  the  funnel 
should  rest  against  the  inside  of  the  receiving  vessel. 

Three  important  factors  determine  the  rate  and  de- 
gree of  separation  by  filtration,  namely:  temperature, 
pressure,  and  the  ratio  between  the  size  of  the  pores  of 
the  partition  and  the  size  of  the  particles  of  the  residue. 

Increase  of  temperature  decreases  the  adhesion  be- 
tween the  molecules  of  the  filtrate  and  those  of  the 
residue,  and  increases  the  size  of  the  colloidal  granules 
and  crystals  of  the  residue. 

Pressure  needful  for  filtration  is  usually  obtained 
through  gravity,  but  sometimes  through  gravity  and  suc- 
tion combined.  The  most  effective  method  of  diminish- 
ing the  atmospheric  pressure  is  the  use  of  the  suction 
pump.  A  platinum  cone  should  be  placed  in  the  apex 
of  the  funnel  to  support  the  moistened  paper,  which 
should  be  so  closely  fitted  to  the  sides  of  the  funnel 
as  to  leave  no  air  channels.  The  funnel  tube  is  then 
to  be  connected  with  the  filter  flask  of  a  suction  pump. 


METHODS   OF  ANALYTICAL   SEPARATION        29 

The  third  factor  for  effective  filtration  consists  in 
increasing  the  size  of  the  particles  of  residue,  or  in 
diminishing  the  size  of  the  pores  of  the  partition.  The 
particles  of  the  residue  are  enlarged  by  heat  and  by 
contact  with  the  liquid  from  which  they  are  formed. 
Both  colloidal  and  crystalline  particles  grow  when  im- 
mersed in  a  mother  liquor  containing  smaller  particles 
of  the  same  kind.  Hence  it  is  the  usual  practice  to 
digest  the  mixture  for  a  short  time  before  filtration. 

Care  should  be  taken,  however,  to  keep  the  residue 
covered  with  liquid  during  the  digestion,  as  on  exposure 
to  the  air  many  precipitates  oxidize  and  often  redissolve 
in  another  form.  The  amorphous  variety  of  ferrous 
sulphide  obtained  by  precipitation  with  ammonium 
sulphide  oxidizes  in  contact  with  the  air  to  soluble 
ferrous  sulphate. 

Experiment  6 

Arrange  filtering  apparatus  as  described  and  filter  the  follow- 
ing mixtures :  silver  chloride  in  water,  barium  sulphate  in  water 
(made  by  adding  dilute  sulphuric  acid,  drop  by  drop,  to  a  dilute 
solution  of  barium  chloride),  and  aluminum  hydroxide  in  water 
(made  by  adding  a  very  dilute  solution  of  sodium  hydroxide 
to  a  concentrated  solution  of  alum  till  a  heavy  gelatinous  mass 
appears).  Try  all  three  mixtures  with  folded  filters,  then  with 
creased  filters,  and  finally  with  the  filter  pump.  After  all  these 
trials,  if  the  filtration  of  either  mixture  is  very  slow  or  the  filtrate 
remains  muddy,  try  decantation  first,  and  then  filtering  the  moist 
residue  with  the  pump. 

Washing.  —  Neither  decantation  nor  filtration  will 
thoroughly  cleanse  the  residue  from  the  filtrate.  It  is 
necessary  in  most  cases  to  wash  off  the  adhering  filtrate 


30  CHEMICAL  ANALYSIS 


^ 


by  pouring  on  water,  or  water  made  acid,  alkaline,  or 
salt,  according  to  the  nature  of  the  residue.  When  a 
thorough  separation  is  demanded,  the  washing  should 
continue  till  the  washings  no  longer  show  a  trace  of  the 
solutes  of  the  filtrate.  In  many  separations,  where  the 
residues  are  liable  to  oxidation  on  exposure  to  the  air, 
it  is  necessary  to  hasten  the  washing  with  the  pump. 

Two  rather  serious  difficulties  are  frequently  en- 
countered both  in  washing  and  in  filtration ;  namely, 
the  clogging  of  the  paper  with  certain  finely  divided 
residues,  and  the  tendency  of  such  residues  to  pass 
through  the  paper.  As  both  difficulties  are  due  to  the 
same  cause,  the  same  treatment  will  correct  both.  The 
troublesome  residues  are  colloidal  in  nature  and,  as  has 
been  stated,  their  particles  unite  or  coagulate  on  heating. 
Hence  it  is  well  to  subject  the  mixture  to  quiet  heat  — 
not  boiling  —  long  enough  to  allow  the  precipitate  to 
settle  to  the  bottom;  decant  the  supernatant  liquid  into 
a  filter  and  apply  the  pump;  add  warm  water  to  the 
residue;  and,  after  settling,  decant  and  filter  again  in 
like  manner.  Repeat  .the  washing  by  decantation  once 
or  twice,  and  then  add  the  solid  to  the  paper,  finally 
washing  directly  on  the  paper. 

As  colloidal  substances  are  somewhat  soluble  in 
water,  but  less  soluble  in  many  neutral  salt  solutions, 
the  latter  are  frequently  used  for  washing.  It  is  also 
necessary  in  special  cases  to  use  acid  or  akaline  solu- 
tions for  washing. 

As  regards  the  effectiveness  of  washing,  Ostwald 
gives  this  calculation :  "  Should  the  washing  liquid 
amount  to  nine  times  as  much  as  the  original  moisten- 


METHODS   OF  ANALYTICAL   SEPARATION        31 

ing  solution,  and  should  1  gram  of  foreign  substance 
be  mixed  with  the  precipitate  to  begin  with,  then  after 
four  washings  only  (J^)*  gram,  i.e.  0.0001  gram,  of  the 
impurity  would  remain." 

Experiment  7 

(a)  Wash  the  residue  of  barium  sulphate  (last  experiment) 
with  water  till  the  washings  show  no  white  precipitate  with  drops 
of  barium  chloride  solution. 

(b)  Precipitate  some  ferric  hydroxide  by  adding  ammonium 
hydroxide  to  a  boiling  solution  of  ferric  chloride  till  the  latter  is 
permanently  alkaline.  Filter  and  wash  according  to  directions 
above. 

PHYSICAL   AND    CHEMICAL   SEPARATIONS 

The  second  class  of  separations,  to  which  reference 
was  made  on  p.  26,  includes  solution,  precipitation, 
evaporation,  and  ignition. 

Solution.  —  Most  methods  of  separation  are  more  or 
less  dependent  upon  solution  as  their  starting  point, 
and  in  some  cases  separation  is  completed  by  this  oper- 
ation. Solids  composed  of  two  or  more  substances,  like 
minerals  and  alloys,  may  be  separated  in  this  manner 
when  only  a  part  of  their  constituents  is  soluble. 

Experiment  8 

(a)  Finely  powder  a  small  piece  of  dolomite  in  an  agate 
mortar,  and  dissolve  by  warming  with  dilute  hydrochloric  acid. 
The  small  residue  insoluble  in  the  acid  is  silica,  whose  separation 
may  be  completed  by  filtration. 

(b)  Dissolve  some  filings  of  soft  solder  in  a  test-tube,  in  a 
mixture  of  equal  parts  of  nitric  acid  and  water,  and  when  the 


32  CHEMICAL  ANALYSIS 

action  ceases,  add  some  water.  When  the  white  powder  has 
settled,  pour  oft"  the  clear  liquid  into  a  small  dish  and  evaporate 
to  dryness.  Try  to  dissolve  the  residue  in  the  test-tube  by  boiling 
with'  water.  Soft  solder  is  an  alloy  of  tin  and  lead,  and  by  solu- 
tion of  the  lead  in  nitric  acid  the  two  metals  are  separated. 

The  ordinary  solvents  for  solids  are  water,  hydro- 
chloric acid,  nitric  acid,  and  aqua  regia. 

A  small  portion  of  the  powdered  solid  is  treated 
in  a  test-tube  with  cold  water.  If  the  solid  does  not 
disappear  after  shaking  the  contents  several  times  at 
intervals,  transfer  a  drop  of  the  solution  to  a  watch 
glass  with  a  glass  rod.  Place  also  a  drop  of  distilled 
water  on  the  watch  glass  near  that  of  the  solution., 
Heat  the  watch  glass  on  an  asbestos  board  or  sand  bath 
to  dryness.  Compare  the  residues  left  from  the  two 
drops.  If  both  are  alike  in  size,  cold  water  does  not 
dissolve  the  solid;  but  if  the  residue  from  the  supposed 
solution  is  larger,  the  solid  is  at  least  partly  soluble  in 
water.  In  either  case  boil  the  contents  of  the  test-tube, 
and  if  the  solid  still  does  not  disappear,  again  evaporate 
a  drop  of  the  solution  to  dryness  on  the  watch  glass.  A 
large  residue  indicates  a  partial  solubility  of  the  solid  in 
hot  water. 

Treat  another  portion  of  the  powdered  solid  in  a 
similar  manner  in  succession  with  dilute  and  concen- 
trated hydrochloric  acid,  dilute  and  concentrated  nitric 
acid,  and  aqua  regia.  If  the  substance  dissolves  com- 
pletely in  any  one  of  the  solvents,  the  solvents  following 
need  not  be  used. 

Sometimes  a  solid  is  composed  of  different  substances 
which  have  no  solvent  in  common.     One  substance  may 


METHODS   OF  ANALYTICAL   SEPARATION        33 

dissolve  only  in  water,  another  only  in  hydrochloric 
acid,  another  only  in  nitric  acid,  etc.  In  such  extreme 
cases  it  is  necessary  to  separate  the  mixed  solids  by 
means  of  solvents.  Some  substances  cannot  be  dis- 
solved by  any  of  the  reagents  mentioned.  It  is  neces- 
sary in  such  cases  to  fuse  them  with  an  alkaline 
carbonate  in  a  platinum  crucible  or  foil,  and  afterwards 
to  digest  the  fused  mass  with  water.  (See  directions 
for  fusion,  p.  43.) 

Experiment  9 

(a)  Using  the  methods  described  above,  dissolve  .5  gram  of 
each  of  the  following  substances :  cupric  sulphate,  barium  sul- 
phate, and  sand. 

(b)  Mix  .5  gram  each  of  the  same  substances  and  separate 
them  by  solution. 

Precipitation. — The  terms  soluble  and  insoluble,  as  used 
in  practice,  have  only  relative  values ;  and  they  merely 
indicate  considerable  differences  in  degree  of  solubility, 
for,  in  fact,  all  substances  are  soluble.  We  have  seen 
in  Exp.  1  that  a  substance  which  is  in  complete  solu- 
tion at  one  temperature  may  be  rendered  less  soluble 
and  thrown  out  of  solution  by  change  to  another  tem- 
perature. And  we  also  have  seen  that  a  material  whose 
solubility  appears  to  be  about  the  same  at  all  tempera- 
tures may  be  thrown  out  of  solution  by  modifying  the 
solvent  in  such  a  manner  as  to  decrease  its  solubility 
therein.  When  by  either  of  these  means  we  have 
forced  a  solution  to  give  up  a  part  of  its  solute  in  solid 
form,  we  have  performed  the  operation  of  precipitation ; 
and  the  solid  thrown  down  is  called  a  precipitate.     In 


34 


CHEMICAL  ANALYSIS 


Acetate. 

Arsenate. 

Arsenite. 

Borate. 

Bromide. 

Carbonate. 

Chlorate. 

Chloride. 

Chromate. 

Cyanide. 

Ferrocyanide. 

Ferricyanide. 

Fluoride. 

Hydroxide. 

Iodide. 

Nitrate. 

Oxalate. 

Oxide. 

Phosphate. 

Silicate. 

Sulphate. 
Sulphide. 
Sulphite. 
Tartrate. 

Ag 

rH(N(M(Neo(NT-ico(NeocoeciH         c<3 

tH    (N    IM    IM 

Tjt    <N    C^    <N 

Ag 

Hg' 

'*(N<N<N<»(NrHa>(M                                         <N 

r-l    (M    (N    (N 

^  (N  e»  tj< 

Hg' 

Pb 

r-l(M(M(M»0(NiHiOO(M(MT^(N(N'* 

iH    (N    (N    N    C^ 

O    (M    C^    <N 

Pb 

Hg"    [rH(N<N             lH(MrtrHTj(rH                      ^             (N 

rl    <N    C»    <M 

tH    (N    «    <N 

Hg" 

Bi|r-.(N        (MTJ^(N1-lTJ<|^^                   ^(N(n 

tH    (N    <N    (N 

rH    <N    <N    <M 

Bi 

Cu|'-l(N<MC^TH(NiHrHrH(MM           (M<Mr-l 

tH    (N    (><    <N    <M 

11    <N    tH    rH 

Cu 

Cd|^(N            Tj<TH(Mi-trH(MCq                    -#(N.-I 

11    (N    (M    <N    <N 

rH    <M    tH    T)< 

Cd 

Sn|i-i              (N              i-HrH(N         eoeorH(N^ 

<N    <N    <N 

r-l    C^    tH    (N 

Sn 

Sb 

(NC^Tt<                   -^IM                          rH<NTH 

C^    (M    •* 

(M    (N    (N    tH 

Sb 

As 



1H             CO    (N 

(M 

A. 

Fe" 

r-i(M(M(NrH<N,H^           «O«e0Tj((NTH 

rH    (M    (N    (N    (N 

y-t    0^    T-t    rfl 

Pe- 

Fe'" 

i-H<N<N(MtH             T-lr-1,-1             eOrtrHClrH 

iH    C^    (M    <M    (N 

rH    (N    11    iH 

Fe- 

Cr 

tH(M              (N»0(NrHlO(M(N                        rH(MrH 

rH    •^    «0    <M    <N 

ij*    «D    rt    11 

or 

Al 

tH(M              (M^              T-t     Tl                                           TH(NrH 

i-H    (N    <M    <N    C^ 

11    <M    (M    ^ 

Al 

Co 

TH(N<N<Ni-l<Nr-lrH(M«OeoeO-*(MiH 

tH    (M    (N    <N    <M 

H    (M    (M    11 

Co 

Ni 

iHC^«(Nf-i(NiHiHc^a>eoeoTj<c<iTH 

rt    <N    <M    (N    (N 

Tl    <N    <N    <N 

,Ni 

Mn 

rH(M(MClTHC»rtiH,H(NC^eO(N(M.-l 

rH    •*    d    (M    (N 

Tl    C<    C^    Tt< 

Mn 

Zn 

tHiM             <N^C<li-(^rH<N«>(N^(NiH 

tH    (N    (M    <M    <N 

n    (N    tH    (N 

Zn 

Ba 

r-l(^^<^^(^^rH(N»Hr-l(^J■*TJ^        o,Hr-i 

^    (N    rt    <N    <N 

CC    n    (M    <N 

Ba 

Sr 

iH(N(M(MTH(Mr-liHTjHT-(rH             (;£,i-HTrl 

iH    (M    iH    (M    <N 

lo  Ti  o  (M  [  Sr 

Ca 

i-IO^<M(MT-ICqrHi-l'*<THi-HrtO'i<TH 

1-1    IN    •*    <N    <N 

lo  11  11  <M  1  Ca 

Mg 

tH(N<N'*i-l<Nr-llH,HTHrH,-lCO(MTH 

tH   ^   d   (M   e< 

^  (N  n  <N  1  Mg 

Na 

»4^^^^^^-^^-^^-,- 

Na 

^^^^^^^^^^^^^^^ 

11    H    rH    r1 

K 

_._^^^^^---^--^^ 

K 

^^^^^^^^^^rHrH^rH^ 

,_!     H     11     tH 

Li 

_^^^^^^^rH^^^,H,-,^ 

Li 

^^^^^^^^^^^^^^^ 

rH     i1     rH     H 

NH, 

NH, 

®     a.                _        .       .                                ^        _ 

Acetate  .    . 
Arsenate     . 
Arsenite      . 
Borate     .    . 
Bromide 
Carbonate  . 
Chlorate 
Chloride  . 
Chromate 
Cyanide  . 
Ferrocyanid 
Ferricyanid 
Fluoride  . 
Hydroxide 
Iodide      . 

Nitrate    . 
Oxalate  . 
Oxide.    . 
Phosphate 
Silicate    . 

Sulphate 
Sulphide 
Sulphite 
Tartrate 

METHODS   OF  ANALYTICAL  SEPARATION        85 

the  two  cases  cited  the  separation  is  of  physical  nature ; 
for,  apart  from  a  decrease  in  the  degree  of  dissociation, 
no  chemical  change  has  been  worked  upon  the  solute. 
There  are  cases,  however,  which  are  more  common 
in  the  practice  of  analytical  operations,  in  which  the 
physical  separation  is  accompanied  with  or  preceded 
by  chemical  changes  without  which  the  precipitation 
would  not  occur.  Such  cases  involve  reaction  between 
the  solute  and  some  added  material,  called  the  precipi- 
tant; and  the  principle  governing  such  reaction  is  that 
the  ions  representing  that  part  of  the  solute  whose 
separation  is  desired,  combine  with  certain  ions  of  the 
precipitant  to  form  a  new  compound  which  is  less  dis- 
sociated and  less  soluble  than  either  the  solute  or  pre- 
cipitant. So  soon  as  enough  of  the  new  and  less  soluble 
compound  has  been  formed  to  saturate  the  solution, 
the  excess  which  is  formed  is  obliged  to  separate  in  the 
form  of  a  precipitate. 

Precipitation  is,  therefore,  the  result  of  supersatura- 
tion,  brought  about  either  by  chemical  or  physical 
means;  but  it  does  not  always  take  place  immediately 
upon  the  occurrence  of  supersaturation.  Frequently 
we  are  obliged  to  resort  to  various  devices  —  such  as 
heating,  cooling,  shaking,  or  stirring — to  induce  it  to 
begin. 

Theories  of  Precipitation. — The  early  chemists  believed 
that  chemical  reactions  are  governed  solely  by  chemical 
affinity.  Bergman  (1775)  taught  that  when  two  com- 
pounds, MX  and  NY^  are  brought  together,  if  the 
affinity  of  Mis  greater  for  Fthan  for  X,  there  will  be 
a  complete  metathesis:   MX+  NY  =  MY  +  NX.     For 


36  CHEMICAL  ANALYSIS 

example:  if  solutions  of  barium  chloride  and  sodium 
sulphate  were  brought  together  in  suitable  proportions, 
there  should  be  complete  precipitation  of  the  insoluble 
barium  sulphate,  because  of  the  greater  affinity  of 
barium  for  sulphuric  acid  than  for  hydrochloric  acid. 
According  to  Bergman's  theory,  reaction  in  the  reverse 
direction  should  not  take  place;  that  is,  barium  sul- 
phate should  not  be  decomposed  by  treatment  in  the 
presence  of  water  with  some  substance  like  sodium  car- 
bonate, for  whose  acid  constitutent  barium  had  less 
affinity  than  for  sulphuric  acid,  with  which  it  already 
was  united. 

BerthoUet  (1804)  afterwards  demonstrated  that  Berg- 
man's theory  is  true  only  in  part,  and  that  there  is  a 
metathesis  between  barium  sulphate  and  sodium  car- 
bonate, provided  a  large  excess  of  the  latter  is  used. 
He  announced  the  theory  that  chemical  reactions  de- 
pend upon  two  things,  —  relative  chemical  affinity 
and  relative  mass,  —  and  embodied  the  theory  in  the 
following  laws,  which  bear  his  name :  — 

1.  When  solutions  of  different  substances  are  mixed, 
and  a  substance  volatile  under  existing  circumstances  can 
be  formed^  it  is  formed  and  escapes, 

2.  When  solutions  of  different  substances  are  mixed, 
and  a  substance  insoluble  under  existing  circumstances 
can  be  formed,  it  is  formed  and  separates. 

The  modern  theory  of  dissociation  enables  us  to  form 
a  clearer  conception  of  these  precipitation  phenomena 
which  result  from  chemical  action.  As  was  said  on 
p.  19,  the  formation  of  an  insoluble  compound  is 
due  to  the  union  of  ions,  from  the  considerably  disso- 


I 


METHODS   OF  ANALYTICAL  SEPARATION        37 

ciated  solute  and  precipitant  to  a  body  whose  dissocia- 
tion and  solubility  are  relatively  less. 

In  the  case  of  barium  chloride  and  sodium  sulphate 
we  have  in  solution,  before  the  materials  are  mixed, 
molecules  of  BaCl2  in  equilibrium  with  ions  of  Ba  and 
CI,  and  molecules  of  NagSO^  in  equilibrium  with  ions 
of  Na  and  SO^.  Upon  mixing,  the  Ba  and  SO^  ions 
are  given  the  opportunity  for  combination. 

Now  the  solubility  of  BaSO^  is  very  slight;  i.e.,  the 
concentration  of  its  solution,  when  the  solute  is  in 
equilibrium  with  solid  BaSO^,  is  very  low.  Yet  this 
concentration  is  made  up  of  the  values  c,  the  concen- 
tration of  the  undissociated  portion,  and  of  a  and  ft, 
the  concentrations  of  the  Ba  and  SO^  ions;^  and  we 
have  seen  that  a.b  =  k.e.  It  will  be  understood  that 
no  more  Ba  and  SO4  ions  can  simultaneously  exist  in 
one  solution  than  will  satisfy  this  equation ;  and  there- 
fore these  ions,  whose  concentrations  in  the  solutions 
of  barium  chloride  and  of  sodium  sulphate  were  very 
considerable,  must  unite  until  the  equation  is  satisfied; 
and  as  the  sum  of  the  concentrations  c  and  h  and  a  is 
very  small,  solid  barium  sulphate  will  be  separated  until 
equilibrium  has  been  restored.  But  as  fast  as  the  ions 
of  Ba  and  SO4  are  thus  withdrawn  from  solution 
further  portions  of  BaClg  and  NagSO^  are  dissociated; 
and  so  the  reaction  continues  until  either  all  of  the 
Ba  or  all  of  the  SO4  in  excess  of  that  permitted  by 
a.b  =  k.e  has  been  used  up. 

The  solubility  of  the  barium  sulphate  is  lessened  by 
the  presence  of  an  excess  of  either  a  soluble  salt  of 

1  For  discussion  of  this  relation,  refer  to  pp.  18-20. 


38  CHEMICAL  ANALYSIS 

barium  or  of  sulphuric  acid.  An  increase  in  the  con- 
centration of  either  the  Ba  or  the  SO4  ions  will  reduce 
the  proportion  of  the  dissociated  salt,  and  consequently 
the  amount  in  solution. 

Since  an  excess  of  reagent  often  vitiates  the  reaction, 
it  is  advisable  to  add  the  precipitant  solution  slowly,  so 
as  to  be  able  to  follow  the  course  of  the  reaction  and 
to  avoid  error. 

Experiment  10 

(a)  In  two  test-tubes  put  a  solution  of  silver  nitrate.  To  the 
first  add  very  little  dilute  ammonia,  drop  by  drop,  shaking  the 
test-tube  after  each  drop.  To  the  second  add  ammonia  directly 
without  dropping. 

(b)  In  each  of  the  test-tubes  put  a  very  dilute  solution  (1  :  500) 
of  silver  nitrate.  To  the  first  add  concentrated  hydrochloric  acid 
and  boil.  When  the  white  precipitate  disappears  add  water. 
To  the  second  add  dilute  hydrochloric  acid  and  boil. 

Evaporation.  —  Evaporation  is  the  removal  of  part  or 
the  whole  of  a  liquid  by  volatilization.  It  is  sometimes 
accomplished  by  exposure  to  the  open  air  or  to  the  sun 
in  flat  vessels  ;  but  for  analytical  purposes  this  pro- 
cedure is  too  slow,  and  application  of  artificial  heat  is 
necessary.  When  very  rapid  evaporation  is  desired 
and  proper  care  is  exercised  to  remove  the  vessel  when 
the  residue  is  quite  dry,  the  operation  may  be  carried 
out  in  an  open  dish  on  a  sand  bath,  or  on  an  iron  or 
asbestos  plate,  over  the  direct  flame.  The  dangers  of 
spattering  and  of  overheating  the  residue  have  put  this 
method  somewhat  in  disrepute.  The  safest  and  most 
satisfactory  method  in  all  respects  is  the  use  of  liquid 
baths.     Water  or  steam  baths  are  most  frequently  used ; 


METHODS   OF  ANALYTICAL  SEPARATION        39 

but  for  evaporations  requiring  high  temperatures,  oil 
and  paraffin  baths  are  sometimes  employed.  In  labora- 
tories supplied  with  water  pressure  and  fixtures,  water 
baths  should  be  supplied  with  constant  level  apparatus 
to  prevent  drying.  All  evaporations  of  active  reagents 
should  be  conducted  under  a  hood  or  out  of  the 
laboratory. 

As  a  means  of  analytical  separation,  evaporation  may 
be  partial  only,  as  when  a  dilute  solution  is  to  be  con- 
centrated; or  it  may  be  complete,  as  when  a  solid  solute 
is  to  be  recovered  from  its  solution. 

Experiment  11 

(a)  Make  a  weak  solution  (1  :  100)  of  cupric  sulphate.  Evapo- 
rate on  a  water  bath  down  to  one-fourth  and  allow  it  to  cool. 

(6)  Repeat  (a)  with  a  similar  solution  of  calcium  chloride. 
If  crystals  fail  to  appear,  evaporate  to  dryness  on  an  asbestos 
plate. 

Ignition.  —  This  process  consists  in  the  application  of 
intense  heat  to  solid  bodies,  and  has  for  its  purposes : 

1.  The  separation  of  gases  from  solids,  as  when  cal- 
cium carbonate  is  converted  into  calcium  oxide  through 
the  removal  of  carbon  dioxide. 

2.  The  separation  of  liquids  from  solids,  as  when 
salts  are  ignited  for  the  removal  of  their  water  of 
crystallization. 

3.  The  separation  of  solids  from  solids,  as  when  the 
readily  volatile  ammonium  compounds  are  driven  off 
from  less  volatile  salts. 

The  Bunsen  Lamp.  —  The  sources  of  heat  for  igni- 
tion in  the  analytical  laboratory  are  the  ordinary  gas 


40  CHEMICAL  ANALYSIS 

lamps  —  usually  Bunsen's  type  —  and  the  blast  lamp. 
By  means  of  a  perforated  thimble  and  a  stopcock  both 
the  air  and  gas  supplies  can  be  regulated  in  a  Bunsen 
burner.  An  excess  of  gas  gives  a  luminous  flame  rich 
in  unconsumed  carbon,  while  an  excess  of  air  produces 
a  non-luminous  blue  flame. 

The  flame  is  composed  of  three  zones  —  an  outer 
zone  of  hot  oxidized  gases  mixed  with  oxygen,  a  medial 
zone  of  hot  reduced  gases  mixed  with  unconsumed 
carbon,  and  an  inner  dark  zone  of  cool  unconsumed 
gases.  The  high  temperature  and  the  free  oxygen 
render  the  outer  zone  favorable  to  processes  of  oxida- 
tion, and  it  is  therefore  called  the  oxidizing  flame. 
The  medial  zone,  with  its  highly  heated  unconsumed 
carbon,  is  the  reducing  flame.  The  hottest  part  of 
the  flame  is  the  junction  of  the  reducing  and  oxidizing 
zones,  where  oxidation  is  most  vigorous. 

The  Blowpipe.  —  For  convenience  in  application  of 
the  different  parts  of  the  flame  a  blowpipe  is  often 
used.  The  type  of  this  instrument  is  a  small  tube 
about  20  cm.  long  with  a  mouthpiece  at  one  end  and 
a  nozzle  tip  of  brass  or  platinum  at  the  other  end.  The 
nozzle  end  is  usually  bent  perpendicular  to  the  shaft 
of  the  blowpipe.  A  constant  blast  of  air  is  fed  into 
the  flame  by  using  the  cheeks  as  a  bellows,  without 
interfering  with  the  regular  breathing.  If  the  cheeks 
are  distended,  air  can  be  forced  out  of  the  mouth  by 
their  contraction  at  the  same  time  when  inspiration  is 
being  effected  through  the  nose.  To  acquire  the  proper 
and  efficient  use  of  the  blowpipe  some  practice  is 
necessary. 


METHODS   OF  ANALYTICAL  SEPARATION        41 

To  obtain  the  best  results  the  air  holes  at  the  base 
of  the  Bunsen  burner  are  closed  so  as  to  give  a  small 
luminous  flame.  It  is  best  to  insert  into  the  burner 
tube  a  smaller  and  somewhat  longer  tube,  compressed 
and  beveled  at  the  top.  This  flattens  the  flame  ;  and 
the  oblique  edge  is  convenient  to  the  operator,  since  the 
flame  is  commonly  directed  downwards.  For  oxida- 
tion the  tip  of  the  blowpipe  is  placed  just  inside  the 
flame ;  and  with  a  strong  blast  the  extreme  tip  of  the 
flame  is  made  to  impinge  on  the  object  to  be  oxidized. 
For  reduction  the  tip  is  placed  on  the  edge  of  the 
flame  and,  using  a  moderate  blast,  the  center  of  the 
flame  is  allowed  to  play  upon  the  object  to  be  reduced. 

Experiment  12 

(a)  Heat  a  small  piece  of  metallic  lead  on  a  piece  of  charcoal 
with  the  oxidizing  flame.  Notice  the  yellow  film  of  metallic 
oxide  on  the  lead. 

(b)  Heat  some  lead  acetate  on  charcoal  with  the  reducing  flame 
to  quiet  fusion.  When  cold,  pick  out  the  metallic  lead  and  test 
its  malleability  on  the  anvil. 

Blast  Lamp.  —  When  higher  temperatures  are  de- 
sired than  can  be  produced  by  the  Bunsen  lamp  or 
mouth  blowpipe,  the  blast  lamp  is  employed.  It  con- 
sists essentially  of  two  concentric  tubes,  of  which  the 
outer  and  larger  conveys  the  gas,  while  the  inner  and 
smaller  supplies  the  air  which  is  furnished  under  pres- 
sure by  a  bellows  or  water  blast. 

Crucibles.  —  The  crucibles  in  common  use  are  small 
cup  -  shaped    vessels    of    porcelain    or    platinum ;    for 


42  CHEMICAL  ANALYSIS 

special  purposes  they  may  be  made  of  silver,  graphite, 
iron,  etc. 

The  method  of  applying  the  heat  is  to  place  the 
crucible  on  a  platinum  or  pipe-clay  triangle  supported 
on  a  ring  stand.  The  ring  is  elevated  at  first,  so  that 
the  flame  will  not  heat  the  crucible  too  rapidly.  After- 
ward the  ring  is  lowered,  or  the  flame  elevated,  so  as 
to  obtain  the  required  heat. 

For  testing  small  amounts  of  material,  fragments  of 
porcelain,  mica  or  asbestos  plates,  or  small  platinum 
foils  (about  2  cm.  square)  are  used  instead  of  crucibles. 
It  is  necessary  to  observe  the  following  precautions  in 
using  crucibles  :  — 

(a)  Do  not  heat  or  cool  a  porcelain  crucible  too 
quickly,  as  in  either  case  it  is  liable  to  crack. 

(h)  Apply  only  the  outer  non-luminous  flame  to  a 
platinum  crucible ;  never  a  yellow-white  flame,  or  the 
interior  blue  zone  of  the  flame,  as  the  one  contains 
unconsumed  carbon,  and  the  other  acetylene  gas,  which 
is  easily  decomposed  into  hydrogen  and  carbon.  At 
high  temperatures  platinum  unites  with  carbon  to 
form  the  carbide  of  platinum,  which  will  oxidize  later, 
■and  cause  the  metal  to  blister. 

(c)  Always  handle  a  hot  platinum  crucible  with  platinum 
tongs,  and  support  it  on  a  platinum  or  pipe-clay  triangle. 

(d)  Never  ignite  the  following  substances  in  plati- 
num crucibles  :  metals,  salts  of  easily  reducible  metals 
(those  of  lead,  silver,  copper,  arsenic,  etc.),  substances 
evolving  chlorine  or  bromine,  sulphites,  phosphates  in 
presence  of  organic  matter,  and  the  nitrates,  cyanides, 
and   hydroxides    of   the  alkalies  and   alkaline    earths. 


METHODS   OF  ANALYTICAL   SEPARATION        43 

All  of  these  readily  corrode  platinum.  In  case  the 
substance  to  be  analyzed  is  of  unknown  character,  it 
should  be  tested  first  on  a  piece  of  platinum  foil. 

(e)  The  platinum  crucible  can  be  cleaned  by  rubbing 
the  surface  with  moist  sea  sand,  applied  with  the  finger. 
Persistent  stains  are  generally  removed  by  concen- 
trated hydrochloric  or  nitric  acid,  or  by  fusion  with 
borax  or  acid  potassium  sulphate.  If  the  latter  sub- 
stance is  used,  allow  the  fused  mass  to  cool  and  then 
dissolve  in  water.  Do  not  attempt  to  break  out  the 
mass,  as  bending  is  injurious  to  platinum  ware. 

Fusion.  —  In  theory,  fusion  is  identical  with  solution. 
As  has  been  pointed  out,  a  solvent  is  essential  to  the 
chemical  reaction  of  soluble  substances,  in  order  that 
they  may  be  brought  into  intimate  contact ;  and,  simi- 
larly, a  flux  is  essential  to  the  chemical  reaction  of 
substances  subjected  to  fusion.  The  flux  may  be  a 
negative  substance  and  not  enter  into  the  reaction,  but 
merely  act  as  a  solvent ;  e.g.,  fluor  spar  in  metallur- 
gical operations  ;  or  the  same  substances  may  act  both 
as  a  solvent  and  as  a  chemical  reagent;  e.g.,  borax 
fused  with  certain  metallic  oxides.  The  fluxes  most 
frequently  used  in  analytical  operations  are  sodium 
carbonate,  potassium  cyanide,  potassium  nitrate,  borax, 
microcosmic  salt,  and  acid  potassium  sulphate.  Often, 
combinations  of  some  of  these  salts  are  used. 

The  so-called  fusion  mixture  is  composed  of  sodium 
and  potassium  carbonates,  mixed  in  the  proportion  of 
their  molecular  weights.  The  mixed  salts  melt  at  a 
lower  temperature  than  either  of  the  individuals.  For 
fusion  and  oxidation,  a  mixture  of  sodium  carbonate 


44  CHEMICAL  ANALYSIS 

and  potassium  nitrate  is  employed;  for  fusion  and 
reduction,  a  mixture  of  sodium  carbonate  and  potas- 
sium cyanide. 

Experiment  13 

(a)  Mix  intimately  some  powdered  insoluble  manganese  dioxide 
with  twice  its  bulk  of  fusion  mixture,  and  heat  to  quiet  fusion  on 
a  platinum  foil  with  the  oxidizing  flame.  When  cold,  dissolve 
the  gi-een  mass  in  water, 

(&)  In  (a)  substitute  chromic  iron  for  manganese  dioxide,  and 
in  addition  to  the  fusion  mixture  use  an  equal  bulk  of  potassium 
nitrate. 

(c)  In  (a)  substitute  feldspar  or  clay  for  manganese  dioxide, 
and  acid  potassium  sulphate  for  the  fusion  mixture. 

Heating  in  Closed  Tubes.  —  Crucible  ignition  is  gener- 
ally an  oxidation  process,  but  closed  tube  ignition  is 
different  inasmuch  as  the  length  of  the  tube  excludes 
the  oxygen  of  the  air  and  prevents  oxidation.  Thus 
the  closed  tube  is  used  for  reduction.  The  tube  is 
usually  made  of  small,  hard  glass  tubing,  about  .5  cm. 
in  diameter  and  10  cm.  long.  First  cut  the  tube  to  the 
proper  length,  and  then  heat  one  end  of  it  in  the  blow- 
pipe flame  till  it  is  closed  and  well  rounded  off.  While 
red  hot  take  it  out  of  the  flame  arid,  holding  it  vertically 
downward,  blow  steadily  till  a  small  bulb  is  formed. 

For  use,  the  bulb  is  about  half  filled  with  the  sub- 
stance, or  with  the  substance  mixed  with  a  flux  or 
reducing  agent ;  and  it  is  heated  first  with  the  smoky 
flame  to  expel  the  moisture  at  a  low  temperature. 
Often  this  heating  is  sufiicient,  but  if  a  very  high  tem- 
perature is  required,  the  non-luminous  flame  must  be 
applied  and  the  tube  heated  to  the  desired  degree. 


METHODS   OF  ANALYTICAL   SEPARATION        45 

Table  II  —  Behavior  op  Substances  heated  in  a 
Bulb-Tube  ^ 


1 

.  Substance  fuses  and  solidifies  again .     .     . 

Compounds  of  alka- 
lies and  of  alkaline 

2 

.  Substance  does  not  fuse,  but  changes  color : 

earths. 

(a)  Chars  and  evolves  carbon  dioxide    .     . 

Organic  substances. 

(6)  Yellow  while  hot,  white  on  cooling  .     . 

Zinc  oxide. 

(c)  White  while  hot,  brown  on  cooling  .     . 

Bismuth  oxide. 

(d)  White  while  hot,  yellow  on  cooling .     . 

Lead  oxide. 

(e)  Orange  while  hot,  yellow  on  cooling     . 

Tin  oxide. 

3 

.  Substance  gives  off  water 

Water  of  crystalliza- 
tion  or  constitu- 
tion. 

■  Ammonium  salts. 

Arsenic            " 

4 

.  Substance  sublimes 

Antimony        " 
Mercury           " 
Sulphur, 

5 

Substance  gives  off  a  gas : 

(a)  Violet  vapors 

Iodine  and  iodides. 

(6)  Red  fumes  of  nitrogen  oxides      .     .     . 

Nitrates    of    heavy 
metals. 

(c)  Oxygen  (tested  with  glowing  match)     . 

Chlorates  and 

nitrates. 

(d)  Carbon  dioxide  (tested  with  lime  water) 

Carbonates  of 

heavy  metals. 

(e)  Sulphur  dioxide  (detected  by  odor)  .     . 

Sulphur 

compounds. 

(/)  Cyanogen  (odor  of  almonds) .... 

Cyanides. 

Experiment  14 

(a)  Heat  some  zinc  sulphate  in  a  bulb-tube,  using  the  lumi- 
nous flame  first  and  then  increasing  the  heat  to  dull  redness. 
Examine  carefully,  for  all  changes ;  —  whether  water  is  given 
off,  the  color  of  the  salt  while  hot  and  when  cold,  etc. 


46  CHEMICAL  ANALYSIS 

(b)  In  a  bulb-tube  heat  some  lead  nitrate.  Test  the  gas  from 
the  bulb  with  a  glowing  splinter  and  note  the  color,  odor,  etc.,  of 
the  gas.     Also  notice  the  behavior  of  the  solid  substance. 

(c)  Heat  some  sodium  acetate  in  a  bulb  and  notice  the  odor 
of  the  evolved  gas,  and  the  change  in  the  solid. 

(</)  Heat  some  ammonium  carbonate  in  a  bulb.  The  salt 
sublimes  and  collects  on  the  cold  part  of  the  tube. 

Platinum  Wire ;  the  Bead  Tests.  —  The  platinum  wire 
is  used  in  two  sets  of  tests  —  those  for  simple  flame 
coloration  and  those  for  bead  coloration.  Flame  colora- 
tion will  be  discussed  under  another  head. 

Certain  easily  fusible  acid  salts,  such  as  acid  potas- 
sium sulphate  and  borax,  have  the  power  of  displacing 
volatile  acids  and  of  combining  with  metallic  oxides  to 
form  fusible  sulphates,  phosphates,  and  borates. 

It  happens  that  many  metals  impart  distinctive  colors 
to  their  fused  mixtures  with  these  salts,  and  that  they 
may  often  be  identified  by  this  means.  A  platinum 
wire,  provided  with  a  glass  handle  and  looped  at  the 
free  end,  serves  as  a  support  for  a  bead  of  fused  borax 
or  other  fused  acid  flux.  The  loop  of  the  clean  wire  is 
heated  to  redness,  dipped  into  the  powdered  flux,  and 
heated  in  the  non-luminous  flame  till  the  loop  is  filled 
with  a  clear  bead.  A  very  small  piece  of  the  substance 
to  be  analyzed  is  stuck  to  the  soft  bead,  and  is  then 
heated  with  the  blowpipe,  first  with  the  oxidizing,  then 
with  the  reducing,  flame.  It  is  necessary  to  continue 
the  blast  till  the  substance  becomes  thoroughly  incorpo- 
rated in  the  bead.  During  the  operation  the  following 
observations  should  be  made  :  whether  the  substance 
fuses,  whether  the  bead  becomes  transparent  or  remains 


METHODS   OF  ANALYTICAL  SEPARATION 


47 


cloudy,  the  color  of  the  bead,  and  its  behavior  in  the 
oxidizing  and  reducing  flames. 

Experiment  15 

(a)  Make  borax  bead^tests  in  both  the  oxidizing  and  reducing 
flames  with  thin  pastes  of  cobalt  nitrate,  chromium  nitrate,  and 
ferrous  sulphate. 

(b)  Fuse  a  small  bit  of  glass  in  a  bead  of  microcosmic  salt.  The 
floating  silica  constitutes  the  skeleton  head,  a  test  for  silicates. 

Table  III  —  Behavior  of  Substances  fused  in  Borax 
OR  Microcosmic  Salt 


Oxidizing  Flame 

Reducing  Flame 

1 

Blue 

Blue .     .     .     ,     - 

Cobalt  salts 

2. 

8 

Green 

Yellow      .... 

• 

Green     .     . 
Dark  green 
Colorless     . 
Gray .     .     . 

• 

Chromium  salts. 
Iron  salts. 

4. 

5 

Amethyst      .... 
Dark  yellow       .     .     . 
Green  (hot),  blue  (cold) 

Manganese  salts. 
Nickel  salts. 

6. 

Red  or  colorless  . 

Copper  salts. 

Charcoal.  —  Charcoal,  as  a  support  for  ignition,  pos- 
sesses the  following  properties :  it  is  a  reducing  agent 
and  greatly  facilitates  reduction  processes ;  it  is  infusi- 
ble and  has  a  low  conducting  power,  and  hence  forms 
an  ideal  crucible ;  and  since  it  is  porous,  it  absorbs  the 
excess  of  fluxes  and  leaves  the  infusible  substances  on 
the  surface.  It  is  used  in  blowpipe  analysis  as  a  sup- 
port and  reducing  agent  combined.  A  small  conical  pit 
is  bored  with  a  knife  blade  or  iron  forceps  handle  near 
one  end  of  a  piece  of  even  grain,  and  the  substance,  or 


48 


CHEMICAL  ANALYSIS 


mixture  of  the  substance  and  a  flux,  is  placed  in  the  pit 
and  heated  in  the  reducing  flame  with  a  blowpipe. 

Experiment  16 

(a)  Heat  a  small  crystal  of  potassium  nitrate  on  charcoal  with 
the  blowpipe.  It  will  flash  up  and  consume  some  of  the  char- 
coal.    This  kind  of  explosive  combustion  is  called  deflagration. 

(b)  Heat  some  lead  oxide  (litharge)  with  three  times  its  weight 
of  fusion  mixture  and  a  little  potassium  cyanide  on  charcoal  with 
the  blowpipe.  After  fusion  continue  to  heat  till  globules  appear. 
When  cold,  pick  out  the  largest  globule  and  test  its  malleability 
on  the  anvil. 

(c)  Heat  some  zinc  oxide  with  fusion  mixture  on  charcoal  with 
the  blowpipe.  Note  the  color  of  the  flame  and  the  color  of  the 
incrustation  around  the  pit  while  hot  and  when  cold.  Moisten 
the  incrustation  with  a  few  drops  of  a  dilute  solution  of  cobalt 
nitrate  and  heat  again.     Note  the  color  of  the  mass. 

Table  IV  —  Behavior  of  Substances  heated  on 
Charcoal  alone 


1.  Substance  deflagrates 

2.  Substance  fuses  and  is  absorbed  in  charcoal 

3.  Substance  is  infusible  and  incandescent    . 

4.  When  infusible,  as  in  3,  treat  with  cobalt 

nitrate  and  heat : 
(a)  Blue 

(6)  Green 

(c)  Pink 

6.  Substance  leaves  incrustation  : 

(a)  White  with  garlic  odor 

(6)    White  without  odor 

(c)  Yellow  (hot),  white  (cold)  .     .     .     .     . 

(d)  Brown 


Nitrates,  chlorates. 
Alkali  salts. 
Alkali-earth  and 

earth  salts. 


Aluminum  oxide 

and  phosphates. 
Zinc  oxide. 
Magnesium  oxide. 

Arsenic  compounds. 
Ammonium  and 
mercury  com- 
pounds. 
Zinc  oxide. 
Cadmium  oxide. 


METHODS   OF  ANALYTICAL   SEPARATION        49 

Table  V  —  Behavior  of  Substances  heated  on 
Charcoal  with  Flux 


1.  A  metallic  bead  is  left : 

(a)  Soft  and  malleable  and  leaves  yellow  incrusta- 
tion  

(6)  Soft  and  malleable  and  leaves  white  incrusta- 
tion   

(c)  Hard  and  malleable  and  leaves  no  incrustation 

(d)  Hard  and  brittle  and  leaves  white  incrustation 

(e)  Hard  and  brittle  and  leaves  yellow  incrusta- 

tion   

2.  Magnetic  particles 

3.  Red  particles 

4.  When  the  moistened  fused  mass  blackens  a  silver 

coin 


Lead. 

Tin. 

Silver. 

Antimony. 

Bismuth. 

{Iron. 
Nickel. 
Cobalt. 
Copper. 
Sulphur 
compounds. 


CHAPTER   IV 

FLAME   COLORATION  AND   SPECTROSCOPY 

Light  and  Color.  —  Light  is  due  to  vibrations  of  the  ether, 
of  great  velocity,  in  directions  perpendicular  to  the  path 
of  the  light,  and  of  frequencies  which  are  inversely  pro- 
portional to  the  lengths  of  the  ether  waves.  The  vibra- 
tions of  lowest  frequency  and  greatest  length  give  rise 
to  light  of  a  red  color.  With  increasing  frequency,  the 
color  changes  successively  to  orange,  yellow,  green,  blue, 
and  violet,  the  last  tint  being  due  to  the  shortest  vibra- 
tions which  can  produce  any  effect  upon  the  eye.  Ordi- 
nary white  light  —  such  as  is  emitted  by  the  sun,  by  the 
glowing  carbon  of  a  flame  or  electric  lamp,  or  by  the  in- 
candescent mantle  of  the  so-called  ''  Welsbach  burner  " 
—  is  made  up  of  vibrations  of  all  lengths,  the  particular 
quality  of  the  light  being  dependent  on  the  proportions  in 
which  the  vibrations  of  different  lengths  are  mingled. 

All  incandescent  solids  emit  a  white  light  whose  char- 
acter is  dependent  upon  their  temperature  rather  than  on 
their  nature;  and  hence  we  cannot  well  characterize 
solids  by  their  appearance  when  incandescent.  Gases  and 
vapors,  on  the  other  hand,  present  distinctive  colors  when 
heated  to  the  point  of  incandescence,  and  they  may  be 
identified  by  means  of  their  color  characteristics. 

Flame  Colorations. — These  colors  may  be  readily  observed 
in  the  following  manner :   a  platinum  wire,  bent  into  a 

60 


FLAME   COLORATION  AND  SPECTROSCOPY       51 

small  loop  at  one  end,  and  fixed  in  a  handle  of  glass  rod 
at  the  other,  is  dipped  into  a  strong  solution  of  the  material 
whose  vapor  color  is  to  be  investigated,  and  is  then  held 
in  the  non-luminous  flame  of  a  Bunsen  lamp.  According 
to  the  nature  of  the  material,  the  portion  of  tlie  flame 
above  the  wire  may  be  colored  more  or  less  intensely. 
Common  salt,  so  treated,  imparts  a  brilliant  yellow  hue 
to  the  flame ;  all  other  salts  of  sodium  behave  in  similar 
fashion,  and  therefore  this  "flame,"  being  given  by 
no  other  substances,  is  taken  as  being  characteristic  of 
sodium.  Treated  similarly,  potassium  compounds  pro- 
duce a  violet  coloration,  calcium  salts  a  yellowish 
red,  etc. 

Experiment  17 

Guided  by  the  above  description,  confirm  the  following  state- 
ment regarding  flame  colorations : 

Table  VI  —  Simple  Flame  Colorations 


1.  Yellow,  obscured  by  blue  glass  i  . 

2.  Violet,  not  obscured  by  blue  glass 

3.  Carmine-red 

4.  Yellow-red 

5.  Deep  red 

6.  Green 

7.  Blue 


Sodium  salts. 
Potassium  salts. 
Lithium  salts. 
Calcium  salts. 
Strontium  salts, 
r  Barium  salts. 
■I  Copper  salts. 
[  Boric  acid. 
Lead  salts. 
Arsenic 

compounds. 
Copper  chloride. 


1  The  violet  flame  of  potassium  salts  is  so  much  less  luminous  than  the 
sodium  flame  that  the  naked  eye  may  fail  to  detect  it  in  presence  of  the 
latter.     But  whereas  the  sodium  flame  is  mostly  obscured  when  viewed 


52  CHEMICAL  ANALYSIS 

Spectroscopy.  —  The  vibrations  of  which  we  assume 
light  to  be  made  up  follow  one  another  through  the 
ether  in  perfectly  straight  lines  so  long  as  their  paths 
are  not  obstructed.  A  line  of  such  vibrations  we  call 
a  rai/;  a  group  of  parallel  rays,  a  beam.  When  the 
path  of  ray  or  beam  is  obstructed,  a  variety  of  things 
may  happen  according  to  the  nature  of  the  obstruction. 
In  case  the  obstruction  is  pervious  to  the  passage  of 
the  ray,  a  portion  of  the  vibrations  may  be  reflected, 
and  the  remainder  will  pass  on  through  the  obstruction. 
Such  rays  as  have  met  the  surface  of  the  latter  exactly 
at  right  angles  will  pass  straight  on;  but  any  which 
have  met  it  obliquely  may  be  more  or  less  deflected  or 
refracted  from  the  prolongation  of  their  former  path. 
The  degree  of  this  refraction  will  be  governed :  — 

(1)  by  the  relation  between  the  optical  "  densities  " 
of  the  media  through  which  the  ray  is  passed ; 

(2)  by  the  vibration  frequency  or  "  wave  length  "  of 
the  ray. 

When  a  beam  of  white  light  is  passed  obliquely  from 
the  air  into  a  denser  material,  such  as  glass,  there  will 
be  a  dispersion  of  its  rays,  those  of  the  shortest  wave 
lengths  being  bent  most  from  their  original  direction, 
so  that  the  former  beam  of  parallel  rays  will  be  spread 
out  into  a  wedge  of  colored  light,  progressing  from 
red  at  one  edge  to  violet  at  the  other.  When  the 
glass  is  in  the  form  of  a  triangular  prism,  the  disper- 
sion will  be  most  perfect;  and  with  the  help  of  such 

through  a  slide  of  cobalt  glass,  that  of  potassium  is  very  slightly  dimmed. 
Accordingly,  when  sodium  is  present,  it  is  always  necessary  to  examine 
the  flame  for  potassium  with  the  aid  of  the  blue  glass. 


FLAME   COLORATION  AND  SPECTROSCOPY       53 

a  prism,  we  may  study  the  character  of  light  from  any 
source.  Obviously,  if  our  light  is  white,  it  will  be 
dispersed  into  a  continuous  band  or  "  spectrum  " ;  if  it 
is  colored,  we  shall  have  strips  of  color  whose  character 
will  indicate  the  composition  of  the  light  under  exami- 
nation. If  our  beam  is  entirely  made  up  of  vibrations 
of  only  a  few  frequencies,  we  shall  have  a  spectrum 
consisting  of  a  few  bright  lines. 

Spectra  can  also  be  produced  by  gratings  of  numer- 
ous thin  parallel  wires  or  of  fine  parallel  etchings  on 
glass  or  metal.  If  light  from  a  narrow  slit  is  viewed 
through  gratings  parallel  witli  the  slit,  some  of  the 
light  can  be  seen  to  pass  unaffected  while  part  of  it 
produces  a  colored  spectrum  on  each  side  of  the  grat- 
ing.    Spectra  thus  produced  are  due  to  diffraction. 

The  Spectroscope.  —  Two  classes  of  instruments  depend- 
ing, respectively,  on  refraction  and  diffraction  are  used 
for  the  analysis  of  light.  Each  class  has  its  advantages 
and  disadvantages.  The  prism  spectroscope  produces 
brighter  spectra,  as  only  a  small  portion  of  the  light  is 
lost  by  reflection  and  absorption.  In  the  case  of  the 
grating  spectroscope,  some  light  passes  unaffected  be- 
tween the  gratings,  some  is  destroyed  by  interference, 
and  only  the  remainder  is  diffracted. 

For  accuracy  the  grating  spectroscope  is  preferable 
for  two  reasons :  first,  the  dispersions  of  the  rays  in 
grating  spectra  are  directly  proportional,  while  those  of 
the  prism  spectra  are  inversely  proportional  to  the  wave 
lengths ;  and,  second,  the  lengths  of  the  prism  spectra 
are  dependent  on  the  material  of  the  prism,  which 
accounts  for  the  fact  that  no  two  prisms  give  uniform 


54  CHEMICAL  ANALYSIS 

dispersions.  But  the  prism  spectroscope,  though  less 
accurate  for  technical  physical  work,  is  simpler  and 
better  adapted  to  ordinary  chemical  analysis. 

A  simple  form  of  the  prism  spectroscope  consists  of 
a  refracting  prism,  or  set  of  prisms,  and  three  small 
telescopes  mounted  on  a  metal  tripod.  One  of  the  tele- 
scopes, called  the  collimator,  has  at  one  end  a  vertical  slit 
of  adjustable  width.  The  rays  of  light,  having  passed 
through  the  slit  and  been  rendered  parallel  by  the  lens 
of  the  collimator,  are  refracted  and  dispersed  by  the 
prism ;  and  the  resulting  spectrum  is  observed  through 
the  eyepiece  of  the  second  telescope.  The  third  tele- 
scope contains  a  horizontal  millimeter  scale  reduced 
about  one-fifteenth.  Light  from  a  white  flame,  placed 
in  front  of  the  scale  telescope,  passes  through  the  scale 
and  is  reflected  on  that  face  of  the  prism  which  stands 
before  the  eye  telescope,  so  that  both  the  image  from 
the  collimator  and  that  from  the  scale  are  seen  at  the 
same  place. 

Another  form  of  ihis  type  of  instrument  is  the  direct- 
vision  spectroscope,  in  which  one  telescope  is  used  for 
both  the  eyepiece  and  collimator,  the  prisms  being 
placed  within  it.  The  Janssen  direcWision  tele- 
scope,^  made  by  the  Geneva  Society,  contains  also  a 
scale  telescope. 

Kinds  of  Spectra.  —  The  spectroscope  shows  three  kinds 
of  spectra  :  — 

1.  The  Continuous  Spectrum,  produced  by  a  white- 
hot  solid  or  liquid.  Solids  and  liquids  emit  rays  of 
many  wave  lengths  and  thus  give  the  colors  blended 
in  the  order  of  their  wave  lengths.     A  platinum  wire 


FLAMK   COLORATION  AND  SPECTROSCOPY       55 

heated  to  whiteness  in  a  non-luminous  flame  shows  a 
continuous  spectrum  in  the  spectroscope. 

2.  The  Discontinuous  Spectrum,  produced  by  an  in- 
candescent gas  or  vapor.  Gases  and  vapors  emit  rays 
of  few  wave  lengths,  and  '  hence  the  spectrum  shows 
only  certain  bright  lines  or  bands.  A  platinum  wire 
dipped  in  a  paste  of  common  salt  and  hydrochloric  acid, 
and  heated  in  a  non-luminous  flame,  shows,  in  addition 
to  the  continuous  spectrum  of  the  white-hot  wire,  a 
bright  yellow  line  produced  by  the  vapors  of  the  salt. 

3.  The  Absorption  Spectrum,  produced  by  an  incan- 
descent solid  or  liquid  which  is  viewed  through  a 
gaseous  or  liquid  medium.     The  medium  absorbs  the 

.rays  peculiar  to  itself,  and  thus  produces  certain  dark 
lines  in  the  continuous  spectrum.  The  so-called  Frauen- 
hofer's  lines  —  dark  bands  which  cross  the  solar  spec- 
trum —  are  caused  by  the  absorption  of  rays  emitted  from 
the  interior  mass  of  the  sun  in  their  passage  through  an 
exterior  gaseous  envelope.  Similarly,  in  the  same  way, 
white  light  passed  through  a  solution  of  potassium 
permanganate  gives  a  spectrum  deprived  of  the  yellow, 
green,  and  blue  rays,  in  whose  places  are  seen  dark 
bands. 

Flame  Spectra.  —  Those  solids  which  vaporize  at  the 
temperature  of  a  Bunsen  flame  can  easily  be  examined 
in  the  flame.  They  are  the  salts  of  the  alkali  and  the 
alkali-earth  metals.  A  blank  analysis  should  first  be 
made  by  holding  a  clean  platinum  wire  in  the  dark 
flame  about  2  cm.  before  the  collimator  slit.  By  focus- 
ing the  telescopes  and  darkening  the  room,  a  dim 
yellow  line  will  be  observed,  due  to  the  presence  of 


56  CHEMICAL  ANALYSIS 

sodium  compounds  in  the  atmosphere.     I'his  may  be 
expected  in  all  spectra. 

For  convenience,  it  is  customary  to  regulate  the  scale 
so  that  the  left-hand  margin  of  the  yellow  sodium  line 
will  exactly  coincide  with  50.^  For  analysis  a  thin 
paste  of  the  salt  is  supported  by  a  small  loop  on  a 
platinum  wire  and  held  in  the  non-luminous  flame 
before  the  collimator  slit. 

Spark  Spectra.  —  Those  solids  which  are  vaporized  not 
by  a  simple  flame,  but  by  an  electric  spark,  include  a 
large  majority  of  the  elements  and  compounds.  The 
spectra  of  gases  are  also  produced  by  the  electric  spark, 
which  can  be  made  by  an  induction  coil  or  influence 
machine. 

For  ordinary  analyses,  a  good  apparatus  consists  of 
a  3-inch  spark  coil  charged  with  a  storage  battery  of 
three  or  five  cells.  Primary  cells  are  either  too  weak  or 
too  inconstant.  The  voltage  of  the  coil  can  be 
increased  by  passing  the  positive  pole  through 
a  Leyden  jar.  A  support  for  the  substance  to 
be  analyzed,  such  as  is  represented  by  the  cut, 
can  be  made  by  fusing  the  end  of  a  platinum 
wire  4  cm.  long  into  one  end  of  a  small,  thin 
glass  tube  1cm.  in  diameter,  so  that  the  wire 
will  stand  free  on  the  inside  of  the  tube  about 
2  cm.  from  the  closed  end.  Cut  the  tube  so 
that  the  edge  of  the  little  cup  will  stand  about 
1mm.  above  the  end  of  the  wire.  Draw  out  another 
piece  of  thin  tubing  to  a  capillary  about  the  length  of 
the  inside  wire.  A  number  of  these  cups  should  be  made 
and  kept  ready  for  use  in  a. test-tube  of  distilled  water. 


FLAME   COLORATION  AND   SPECTROSCOPY       57 

I  In  using  this  apparatus,  first  connect  the  empty  cup 
I  to  the  negative  pole  of  the  coil  by  means  of  a  small 
I  U-tube  filled  with  mercury.  Clamp  the  positive  plati- 
num tipped  pole  above  the  cup  so  that  the  two  poles 
will  be  about  1mm.  apart.  Close  the  circuit  and  make 
a  blank  analysis  with  the  spectroscope.  Certain  brigjit 
lines  representing  the  spectra  of  the  gases  of  the  air 
are  often  seen.  In  order  to  avoid  error  in  subsequ^t 
analyses  for  spark  spectra,  the  presence  of  thes^  lir^es 
should  be  anticipated.  Fill  the  cup  about  one-thijrd 
full  of  a  strong  solution  of  the  substance  to  be  analyzed, 
close  the  circuit,  and  examine  the  spectrum.  ' 

Absorption   Spectra.  —  Solutions   of   many  substances, 
:   both  inorganic  and  organic,  and  also  many  gases,  give 
characteristic  absorption  spectra. 

The  solvents  are  various,  though  water  and  alcohol 
are  most  common.  This  method  not  only  confirms 
many  line  spectra  of  inorganic  compounds,  but  also 
affords  the  only  means  of  spectrum  analysis  for  com- 
pounds decomposed  by  heat.  Among  the  inorganic 
bodies  whose  absorption  spectra  are  important  to  the 
analyst  are  the  salts  of  aluminum,  iron,  cobalt,  nickel, 
arid  manganese.  The  most  important  among  organic 
compounds  are  the  dyes,  blood,  chlorophyll,  etc.   I       ; 

The  apparatus  necessary  for  producing  absorption 
spectra  is  the  spectroscope,  a  white  light,  and  a  large 
test-tube  to  contain  the  solution  to  be  examined.  First 
arrange  the  spectroscope  as  for  flame  or  spark  spectra, 
and  then  place  the  white  light  about  2  dm.  in  front  ;of 
the  collimator  slit  so  that  a  clear  continuous  spectrum 
will  appear.     Interpose  between  the  slit  and  the  white 


58 


CHEMICAL  ANALYSIS 


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FLAME   COLORATION  AND  SPECTROSCOPY       69 


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60  CHEMICAL  ANALYSIS 


EXPLANATION   OF   TABLE   VII 

Nos.  1,  2,  3,  4,  5,  and  6  represent  the  discontinuous  spectra  of 

;  salts  of  metals  volatile  in  the  flame.     The  lines  and  curves  in  the 

!  field  of  ieach  spectrum  indicate  the  position  and  distinctness  of 

'  visible  lines.     For  example,  the  spectrum  of  potassium  appears 

on  the  scale  as  a  strong  line  between  10  and  20,  —  more  accurately 

:  I7j^  —  and  another  thinner  and  shorter  line  between  150  and  160, 

—  more  accurately  154.     The  long  curve  from  20  to  130  shows 

that  there  are  many  indistinct  lines  within  that  area,  and  the 

varying  heights  of  the  curve  indicate  the  relative  distinctness  of 

the  lines. 

Nos.  7,  8,  and  9  represent  the  discontinuous  spectra  of  salts  of 
metals  volatilized  by  the  electric  spark.  Spark  spectra  are  char- 
acterized by  the  small  number  of  narrow  lines  and  the  absence 
of  indistinct  lines.  No.  10  represents  the  absorption  spectrum  of 
sunlight,  showing  the  so-called  Fraunhofer's  lines. 

Nos.  11,  12,  13,  14,  15,  and  16  represent  absorption  spectra. 
The  shaded  parts  show  the  portion  of  the  spectrum  absorbed,  and 
the  curved  margins  the  relative  degrees  of  absorption. 


FLAME   COLORATION  AND  SPECTROSCOPY       61 

light  the  test-tube  filled  with  a  dilute  solution  of  the 
substance.  Certain  portions  of  the  continuous  spectrum 
will  now  appear  dark. 

Mapping  Spectra.  —  Two  methods  have  been  adopted 
for  recording  spectra :  — 

1.  Kirchoff  and  Bunsen's  scale,^by  which  the  posi- 
tions of  the  lines  are  recorded  on  a  graduated  scale. 
The  conventional  practice  is  to  adjust  the  scale  so  that 
the  yellow  sodium  line  shall  coincide  with  50  on  the 
scale.  This  method  is  quite  simple,  and  though  not  so 
accurate  as  the  other  method,  it  is  generally  used  for 
chemical  analysis. 

2.  The  wave-length  method,  by  which  the  wave 
lengths  of  the  colors  are  calculated  from  the  formula 

V 

X  =  -,  in  which  X,  v,  and  n,  respectively,  are  wave  length, 

velocity  of  light,  and  number  of  vibrations.  The  unit  is 
one  ten-millionth  of  a  millimeter,  called  an  Angstrom. 

Professor  Rowland  of  Johns  Hopkins  University,  by 
means  of  his  improved  concave  grating  spectroscope, 
has  compiled  an  atlas  ^  of  a  large  number  of  spectra 
recorded  in  wave  lengths.  In  this  elementary  book 
measurements  of  wave  lengths  would  not  be  consistent 
with  the  character  of  the  work.  Hence  the  use  of 
Kirchoff  and  Bunsen's  scale  is  recommended. 

The  table  on  pages  58  and  59  includes  some  illustra- 
tions of  a  method  of  mapping  spectra. 

Experiment  18 

(a)  Examine  and  map  the  flame  spectra  of  the  following 
salts:  sodium  chloride,  potassium  chloride,  lithium  chloride, 
barium  chloride,  strontium  chloride,  and  calcium  chloride. 


62  CHEMICAL   ANALYSIS 

(6)  Examine  and  map  the  spark  spectra  of  the  following 
salts:  magnesium  chloride,  zinc  chloride,  manganese  chloride, 
copper  chloride,  and  bismuth  chloride. 

(c)  Examine  and  map  the  absorption  spectra  of  the  follow- 
ing inorganic  salts :  ferric  chloride  in  water,  potassium  perman- 
ganate in  water,  chrome  alum  in  water,  and  cobalt  nitrate  in 
alcohol. 

((/)  Examine  and  map  the  absorption  spectra  of  alcoholic  solu- 
tions of  blood  and  fuchsine,  and  a  water  solution  of  logwood. 

Special  Method  for  Aluminum  (Vogel).  —  Make  a  solution 
of  extract  of  logwood  by  boiling  the  chips  in  water. 
Place  a  test-tube  containing  this  extract  between  the 
spectroscope  and  a  luminous  flame.  The  right  end  of 
the  spectrum  will  be  absorbed,  the  extent  of  absorption 
depending  on  the  concentration  of  the  logwood.  The 
boundary  between  the  absorbed  and  unabsorbed  parts 
of  the  spectrum  is  made  to  coincide  with  a  convenient 
line  on  the  scale.  Now  add  a  few  drops  of  a  dilute 
neutral  solution  of  an  aluminum  salt.  This  will  cause 
the  boundary  line  to  move  to  the  left  in  proportion  to 
the  concentration  of  the  solution.  The  aluminum  salt 
solution  is  made  neutral  by  adding  to  it,  drop  by  drop, 
a  very  dilute  solution  of  ammonia,  until  a  slight  but 
permanent  precipitate  is  produced. 

Neutral  ferric  salts  give  the  same  reaction,  but  iron 
can  be  tested  for  in  the  wet  way.  In  case  of  a  mixture 
of  aluminum  and  iron  salts,  the  iron  can  be  removed 
by  adding  an  excess  of  ammonium  sulphocyanate 
solution  and  shaking  out  the  ferric  sulphocyanate 
with  ether.  The  colorless  aqueous  portion  is  tested 
for  aluminum  salts.  (See  Nos.  13  and  14  on  the 
table.) 


FLAME   COLORATION  AND  SPECTROSCOPY       63 

Special  Method  for  Magnesium.  —  Make  a  dilute  solution 
of  alcana  and  record  its  absorption  spectrum.  Now  add 
a  dilute  neutral  solution  of  magnesium  chloride.  The 
alcana  spectrum  will  be  moved  to  the  left  in  propor- 
tion to  the  concentration  of  the  magnesium-chloride 
solution. 

Special  Method  for  Manganese.  —  Boil  the  compound 
with  some  lead  dioxide  and  a  little  nitric  acid  and 
test  for  the  absorption  spectrum  of  permanganic  acid. 
(See  No.  11  on  the  table.) 

Special  Method  for  Cobalt  (Wolff).  —  Add  ammonium  sul- 
phocyanate  to  the  cobalt-chloride  solution  and  shake 
with  alcohol  (amyl  preferable)  and  ether.  This  dis- 
solves the  cobaltous  sulphocyanate,  and  the  solution 
gives  a  characteristic  absorption  spectrum. 

Special  Method  for  Iodine  (Vogel).  —  Iodine  can  be  tested 
for  with  the  apparatus  here  shown :  a  is  the  gas  flame 
colored  by  the  iodide  ;  6  is  a  hard 
glass  tube  held  in  position  over 
the  smaller  tube  by  a  spiral  cop- 
per wire,  c;  c?  is  a  hard  glass 
test-tube  containing  at  its  both 
tom  a  mixture,  e,  of  copper  oxide 
and  an  iodide ;  /  is  a  stream  of 
illuminating  gas  passing  through 
the  apparatus  and  burning  at  a ; 
g  is  the  spectroscope  ;  A  is  a  gas 
burner. 

Bromine  and  chlorine  can  also  be  detected  in  the 
same  manner  by  using  a  bromide  or  a  chloride  instead  of 
an  iodide.     The  copper  iodide,  or  chloride,  or  bromide 


64  CHEMICAL   ANALYSIS 

escapes  with  the  gas  and  colors  the  flame  green.  The 
spectroscope  shows  a  number  of  bands,  especially  in  the 
green  part  of  the  spectrum,  which  are  different  for 
iodine,  chlorine,  and  bromine. 


CHAPTER   V 

LIST   AND  PREPARATION   OF   REAGENTS 

It  is  desirable  that  the  student  know  not  only  the 
chemical  nature  of  reagents,  but  also  their  proper  dilu- 
tion. From  a  careful  study  of  the  principles  of  solu- 
tion and  of  mass  action,  the  reason  for  a  knowledge  of 
dilution  must  be  obvious. 

In  Exp.  4  an  illustration  is  given  of  the  different 
effects  of  concentrated  and  dilute  sulphuric  acid  on  zinc. 

Furthermore,  it  is  desirable  that  reagents  be  so  diluted 
as  to  give  uniform  strengths,  so  that  the  volume  of  solu- 
tion used  will  be  an  index  of  the  quantity  of  reagent 
employed.  Most  analysts  use  an  arbitrary  system  of 
dilutions  that  has  no  special  significance,  save  that  it 
meets  the  empirical  requirements  of  ordinary  analysis. 

Reddrop  (1890)  suggested  that  dilutions  be  based  on 
the  equivalent  weights  of  the  reagents.  The  equiva- 
lent weight  of  a  substance  is  its  molecular  weight 
divided  by  the  number  of  its  replaceable  hydrogen 
atoms,  or  those  which  are  the  equivalents  of  hydrogen. 
For  example,  the  equivalent  weight  of  hydrochloric  acid 
is  36.5,  obtained  by  dividing  its  molecular  weight  by  its 
number  of  replaceable  hydrogen  atoms : 

36.5^1  =  36.5. 
Likewise,  the  equivalent  weights  of  sodium  chloride, 
sodium  hydroxide,  and  silver  nitrate  are,  respectively, 

65 


66  CHEMICAL  ANALYSIS 

58.5,  40,  and  170.  The  equivalent  weight  of  sul- 
phuric acid  is  49,  obtained  by  dividing  its  molecular 
weight  by  its  number  of  replaceable  hydrogen  atoms : 

98  -^  2  =  49. 

The  following  list  gives  some  equivalent  weights : 

Hydrogen,  H 1. 

Oxygen,  O       8. 

Hydrochloric  Acid,  II CI 36.5. 

Nitric  Acid,  HNO3 63. 

Sulphuric  Acid,  n2S04 49. 

Acetic  Acid,  HC2II3O2 60. 

Phosphoric  Acid,  H3PO4    .......  32.6. 

Ammonium  Chloride,  NH4CI 53.7. 

Barium  Chloride,  BaClz 103.8. 

When  equivalent  weights  are  dissolved  in  equal 
volumes  of  water,  equal  fractional  parts  of  the  solu- 
tions will  be  equivalent.  A  normal  solution  of  a  sub- 
stance is  one  which  contains  the  equivalent  weight  of 
that  substance,  in  grams,  dissolved  in  a  liter  of  solution. 
For  example,  a  normal  solution  of  sodium  chloride  is 
its  equivalent  weight,  58.5  grams,  dissolved  in  sufficient 
water  to  produce  a  liter  of  solution. 

Equal  volumes  of  normal  solutions  whose  solutes 
react  with  one  another  should  neutralize  each  other  per- 
fectly, leaving  no  excess  of  either  reagent,  as  in  the  fol- 
lowing equation :  AgNOg  +  NaCl  =  AgCl  +  NaNOg. 

Experiment  19 

Arrahge  two  burettes.  Fill  the  one  with  a  normal  solution  of 
sulphuric  acid,  and  the  other  with  a  normal  solution  of  sodium 
hydroxide.     Into  a  beaker  or  flask  measure  out  20  c.c.  of  the 


LIST  AND  PREPARATION   OF  REAGENTS         67 

alkali  and  add  about  1  c.c.  litmus  solution.  Now  add  the  acid 
from  the  burette,  drop  by  drop,  stirring  or  shaking  constantly, 
till  the  blue  color  changes  to  red.  Compare  the  volumes  of  the 
two  solutions  used. 

Note.  —  The  normal  solutions  in  this  experiment  should  be  prepared 
by  the  instructor. 

Theoretically,  solutions  of  all  reagents  should  be  of 
equivalent  strengths ;  but  it  has  not  been  found  practi- 
cal to  give  all  of  them  these  values  because  of  their 
lack  of  uniform  solubilities.  It  is  convenient  to  adopt 
the  normal  solution  as  a  standard  and  to  express  all 
dilutions  as  multiples  or  fractions  of  normal. 

The  letter  N  is  used  to  denote  a  normal  solution,  and 
all  variations  from  normal  are  expressed  in  terms  of  N. 
For  example,  the  best  dilution  for  sulphuric  acid  is 
about  five  times  its  normal  strength,  expressed  thus: 
5  N  H2SO4,  or  5  N  solution ;  and  the  best  dilution  for 
silver   nitrate   is    about  one-fifth  its  normal  strength, 

N  N 

expressed  thus :  ^  AgNOg,  or  —  solution. 

It  is  sufficient  for  the  purposes  of  qualitative  analysis 
if  the  strength  of  the  reagents  be  approximately  known, 
and  it  will  be  understood  that  the  strengths  given  below 
are  only  approximate. 

LIST  OF  REAGENTS 

Solutions 1.  Acetic  Acidj  HC2H3O2  +  Aq.     1  volume 

80%  acid  to  2  volumes  water  =  5  N  solution. 

2.  Concentrated  Hydrochloric  Acid,  HCl.  Sp.  gr.  1.20  = 
UN  solution. 

3.  Dilute  Hydrochloric  Acid,  HCl  +  Aq.  1  volume  con- 
centrated acid  to  3  volumes  water  =  5  N  solution. 


68  CHEMICAL  ANALYSIS 

4.  HydrosulphuriG  Acid  Gas,  HgS.  For  generating  the 
gas  Kipp's  apparatus  is  generally  used.^  The  generator 
should  be  kept  in  a  hood  with  a  good  draught.  The  mate- 
rials used  in  the  generator  are  lumps  of  ferrous  sulphide, 
FeS,  and  dilute  hydrochloric  or  sulphuric  acid.  Often 
when  the  acid  seems  exhausted  it  will  renew  its  activity  if 
it  is  removed  from  the  generator,  and  the  lumps  of  ferrous 
sulphide  thoroughly  washed  with  water. 

5.  Hydrosulphuric  Acid  Solution,  HgS  +  Aq.  The  gas  is 
passed  into  cold  water  to  saturation.  The  solution  of  the 
gas  should  be  kept  in  the  dark  or  in  bottles  of  deeply 
colored  glass,  as  sunlight  decomposes  the  acid  with  sepa- 
ration of  sulphur. 

6.  Concentrated  Nitric  Acid,  HNO3.  Sp.  gr.  1.42  =  16  N 
solution. 

7.  Dilute  Nitric  Acid,  HNO3  -f-  Aq.  5  volumes  concen- 
trated acid  to  11  volumes  water  =  5  N  solution. 

8.  Concentrated  Sulphuric  Acid,  H2SO4.  Sp.  gr.  1.84 
=  36  N  solution. 

9.  Dilute  Sulphuric  Acid,  H2S04'  2  H2O  +  Aq.  1  volume 
concentrated  acid  to  6  volumes  water  =  5  N  solution.  In 
diluting  the  concentrated  acid,  it  should  be  added  to  the 
water  very  slowly,  in  a  large  porcelain  dish,  and  allowed  to 
cool  before  using. 

10.  Tartaric  Acid,  H2C4H4OC  +  Aq.  1  part  crystals  to 
13  parts  water  =  N  solution.  The  acid  decomposes  in 
solution  and  should  be  prepared  fresh  each  time. 

11.  Aqua  Regia,  HCl  -h  HNO3.  1  volume  concentrated 
nitric  to  3  volumes  concentrated  hydrochloric  acid.  This 
proportion  is  sometimes  varied  for  specific  purposes.  As  a 
solvent,  just  enough  of  the  reagent  should  be  used  to  dis- 
solve the  substance  completely.  A  large  excess  must  be 
avoided,  since  if  allowed  to  remain  it  decomposes  other 
reagents,  while  if   evaporated  certain  volatile  chlorides, 


LIST  AND  PREPARATION   OF  REAGENTS         69 

e.f/.,  arsenic  and  mercuric  chlorides,  are  liable  to  be  lost. 
Prepare  aqua  regia  fresh  each  time  it  is  needed. 

12.  Ammonium  Chloride,  1^11401 -\-Aq.  Use  1  part  crys- 
tals to  4  parts  water,  and  allow  it  to  rise  to  the  natural  tem- 
perature ;  then  dilute  with  1  part  water  =  5  N  solution. 

13.  Ammonium  Carbonate,  (NH4)2C03  +  Aq  +  NH4OH. 

4  parts  solid  ammonium   carbonate  dissolved  in  7  parts 

5  N  NH4OH ;    then   dilute    with    14   parts   water  =  5  N 
solution. 

14.  Ammonium  ^^Sesqui^^  Carbonate.  Dissolve  1  part 
solid  (NH4)2C03  in  9  parts  water  and  add  for  each  10  c.c. 
of  the  liquid  5  drops  of  strong  ammonia  (Hager).  Pre- 
pare fresh  each  time. 

15.  Concentrated  Ammonium  Hydroxide,  NH4OH  +  Aq. 
Sp.  gr.  0.90  =  9N  solution. 

16.  Dilute  Ammonium  Hydroxide,  NH4OH  -f  Aq.  2  vol- 
umes concentrated  ammonia  water  (sp.  gr.  0.90)  to  5  volumes 
water  =  5  N  solution.  Both  concentrated  and  dilute  ammonia 
attack  glass  vessels  ;  and  if  white  flakes  appear  in  the  clear 
solutions,  they  should  be  filtered  out  before  use. 

17.  Ammonium  Oxalate,  (NH4)2C204- HgO  +  Aq.     1  part 

solid  crystals  to  25  parts  water  =  —  solution. 

18.  Ammonium  Sulphide,  (^114)28  +  Aq  -f-  NH4OH. 
Saturate  3  parts  5N  ammonia  with  hydrogen  sulphide 
gas,  and  then  add  2  parts  5N  ammonia  solution.  This 
reagent  should  be  made  frequently,  as  it  decomposes  on 
standing. 

19.  Yellow  Ammonium  Sulphide,  (NH4)2Sx  -{-  Aq  -\-  NH4 
OH.  Digest  a  solution  of  (NH4)2S  with  a  little  powdered 
roll  sulphur.  An  excess  of  sulphur  must  be  avoided,  as  it 
produces  the  red  solution  containing  higher  sulphides. 

20.  Barium  CA^oric^e,  BaCV  2  H2O  +  Aq.  1  part  solid 
crystals  dissolved  in  10.  parts  water  =  N  solution. 


70  CHEMICAL  ANALYSIS 

21.  Bromine,  Br.  Should  be  kept  in  a  dark,  glass- 
stoppered  bottle. 

22.  Dilute  Bromine,  Br  +  Aq.      Make  a  saturated   so- 

N 
lution  by   shaking  an  excess  of  bromine    in   water  =  — 

solution.     Stronger  solutions  of  bromine  can  be  made  by 
adding  potassium  bromide  to  the  water  solution. 

23.  Calcium  Hydroxide  (lime  water),  Ca(0H)2  +  Aq. 
Saturate  freshly  boiled  distilled  water  by  shaking  an 
excess   of   freshly   slaked   lime   in   it,  and   allowing   ithe 

N 
excess  of  lime  to  settle.     Filtrate  =  —  solution. 

24.  Chlorine    Water,    CI  +  Aq.      Saturate     cold    water 

N 
with  chlorine  gas  =  —  solution.     Should  be  kept  in  the 
o 

dark  and  in  brown  bottle. 

25.  Cobalt  Nitrate,  Co(N03)2-6H20  +  Aq.  1  part  solid 
crystals  to  7  parts  water  =  N  solution. 

26.  Ferric  Chloride,  FeCla  4-  Aq.  1  part  solid  salt  to 
20  parts  water  =  N  solution. 

27.  Ferrous  Sulphate,  FeS04-7H20  +  Aq.  1  part  solid 
crystals  to  7  parts  water  =  N  solution.  It  is  best  to  pre- 
pare this  solution  fresh,  when  needed. 

28.  Lead  Acetate,  Pb(C2H302)2' 3  HgO  +  Aq.  4  parts 
solid  to  21  parts  water  =  N  solution. 

29.  Magnesium  Sulphate,  MgS04-7H20  +  Aq.  1  part 
crystals  to  8  parts  water  =  N  solution. 

30.  Magnesia  Mixture,  MgClg  +  NH4CI  +  ]SrH:40H  +  Aq. 
Dissolve  6  grams  magnesium  chloride  crystals  and  165 
grams   ammonium   chloride  in  300  c.c.  water;  then  add 

300  c.c.  5  N  ammonia  solution ;  and  dilute  to  1  liter  =  — ■ 
solution. 

31.  Mercuric   Chloride,  HgCl2  +  Aq.     1  part  solid  salt 

to  37  parts  water  =  —  solution. 


LIST  AND  PREPARATION  OF  REAGENTS         71 

32.  Hi/drochlorplatinie  Acid,  HoPtClc'G  HgO  +  Aq.  1  part 
solid  salt  to  12  parts  water  =  N  solution.  The  solution 
can  also  be  prepared  by  dissolving  0.30  gr.  platinum  foil 
in  aqua  regia,  evaporating  to  dryness,  and  redissolving  in 
10  c.c.  5  N  HCl. 

33.  Potassium  Chromate,  KaCrOi  -f  Aq.  1  part  solid 
salt  to  10  parts  water  =  N  solution, 

34.  Potassium  Cijanide,  KCN  +  Aq.  1  part  solid  salt 
to  15  parts  water  =  N  solution.  Prepare  fresh  for  each 
experiment. 

35.  Potassium  Ferricyanide,  K3Fe(CN)6-3H20  +  Aq. 
1  part  solid  salts  to  9  parts  water  =  N  solution. 

36.  Potassium  Ferrocyanide,  K4Fe(CN)c*3H20  +  Aq. 
1  part  solid  salt  to  10  parts  water  =  N  solution. 

37.  Potassium  Hydroxide,  KOH  -f  Aq.  1  part  solid 
caustic  potash  to  3.5  parts  water  =  5  N  solution. 

38.  Potassium  Sulphocyanate,  KCNS  +  Aq.  1  part  solid 
salt  to  10  parts  water  =  N  solution. 

39.  Silver  Nitrate,  AgNOa  +  Aq.     1  part  solid  salt  to 

N 

30  parts  water  =  —  solution. 

40.  Sodium  Acetate,  ]SraC2H302-3H20  +  Aq.  1  part  solid 
salt  to  8  parts  water  =  N  solution. 

41.  Sodium  Carbonate,  Na2C03' 1 OH2O -|- Aq.  1  part 
solid  crystals  to  7  parts  water  =  N  solution. 

42.  Sodium  Hydroxide,  NaOH  -f-  Aq.  1  part  solid  to  5 
parts  water  =  5  N  solution.  First  dissolve  the  solid  base 
in  a  little  water,  allow  to  cool,  and  then  dilute  to  the 
required  volume. 

43.  Sodium  Phosphate,  HNa2P04'12H20  +  Aq.  1  part 
solid  crystals  to  8  parts  water  =  N  solution. 

44.  Stannous  Chloride,  SnCl2*2H20  +  Aq.  Dissolve  3 
parts  solid  salt  in  3  parts  5  N  hydrochloric-acid  solution, 
and  dilute  with  20  parts  water  =  N  solution.     Pieces  of 


72  CHEMICAL  ANALYSIS 

granulated  tin  should  be  kept  in  the  solution.  An  excel- 
lent quality  of  solid  stannous  chloride  can  be  made  by 
heating  granulated  tin  with  repeated  small  quantities  of 
concentrated  HCl,  added  at  intervals  whenever  ebullition 
ceases.  When  all  the  tin  is  dissolved,  evaporate  to  dry- 
ness on  the  water  bath. 

Solvents.  —  45.  Alcohol^  CaHgOH,  95  per  cent. 

46.  Carbon  Disulphide,  CSg. 

47.  Ether,  (€2115)20,  commercial. 

48.  Ether-Alcohol,  1  volume  absolute  ether  to  1  volume 
absolute  alcohol. 

49.  Petroleum  Ether. 

50.  Water,  distilled. 

Dry  Reagents.  —  51.  Ammonmm  Chloride,  NH4CI. 

52.  Ammonium  Carbonate,  (NH4)2C03. 

53.  Cobalt  Nitrate,  00(^03)2* 6 H2O. 

54.  Lead  Peroxide,  PbOg. 

55.  Manganese  Peroxide,  MnOg. 

56.  Microcosmic  Salt,  HlSra(NH4)P04-8H20. 

57.  Potassium  Carbonate,  K2CO3,  anhydrous. 

58.  Potassium  Cyanide,  KCN. 

59.  Potassium  Disulphate,  HKSO4. 

60.  Potassium  Nitrate,  KNO3. 

61.  Potassium  Nitrite,  KNOg. 

62.  Sodium  Carbonate,  NagCOg,  anhydrous. 

63.  Sodium  Tetraborate  (borax),  NagBiOy  *  10  HgO. 

64.  Sodium  Peroxide^  Na202. 


CHAPTER  VI 

SYSTEMS   OF  ANALYTICAL  EXAMINATION 

Analysis  by  the  Dry  Way.  —  The  chief  operations  in 
analysis  by  this  system  are  the  observations  of  (a) , 
Oxidation  and  Reduction,  and  (h)  Flame  Coloration. 

(a)  Oxidation  and  reduction  have  been  explained 
under  ignition  operations.  These  include  fusion  in 
crucibles,  closed  tube  reductions,  oxidation  and  reduc- 
tion with  fluxes  on  a  platinum  wire,  and  reduction 
on  charcoal  with  and  without  fluxes. 

(h)  Flame  colorations  have  been  explained  under 
simple  flame  colorations  and  spectroscopy. 

Though  these  operations  are  indispensable  to  the 
analyst,  they  do  not  constitute  an  independent  system, 
but  are  only  used  for  preliminary  and  confirmatory 
observations. 

Analysis  by  the  Wet  Way.  —  Solution  is  the  basis  of 
this  system  of  analysis.  Advantage  is  taken  of  the 
following  facts :  — 

(a)  The  metallic  ions  of  most  compounds  behave 
alike  towards  certain  reagents,  regardless  of  the  acid 
radicals  which  may  be  present.  For  example,  all  solu- 
ble silver  salts  will  give  insoluble  silver  chloride  with 
all  soluble  chlorides. 

(b)  Differences  of  solubility  of  similar  compounds  of 
different  metals  may  be  utilized  in  separating  them  into 

73 


74  CHEMICAL  ANALYSIS 

groups.  For  example,  silver,  lead,  and  copper  sulphides 
are  insoluble  in  acidified  solutions,  while  some  other 
sulphides,  e.g.,  those  of  zinc  and  barium,  are  soluble 
under  like  conditions.  Silver  and  lead  chlorides  are 
insoluble,  and  copper  chloride  is  soluble  in  acidified 
solutions.  Such  differences  of  solubility  afford  an  easy 
means  of  separating  and  detecting  these  metals. 

(c)  Physical  characteristics,  color,  odor,  etc.,  are  used 
for  detecting  individual  substances.  For  example,  the 
soluble  salts  of  both  cadmium  and  copper  are  precipi- 
tated by  hydrogen  sulphide;  but  as  the  one  sulphide 
is  a  bright  yellow  and  the  other  black,  the  two  can  be 
distinguished. 

(d)  The  principles  of  the  periodic  law  of  elements 
are  in  part  regarded  in  analytical  chemistry.  The 
periodic  groups  —  sodium,  potassium,  lithium ;  barium, 
strontium,  calcium ;  chlorine,  bromine,  and  iodine  — 
are  also  utilized  as  analytical  groups. 

An  ideal  natural  system  of  classification  would  have 
analytical  groups  to  coincide  with  periodic  groups ;  but 
in  this  respect  analytical  classification  is  somewhat  arti- 
ficial, and  depends  more  upon  differences  in  degree  of 
the  physical  property  of  solubility  than  on  chemical 
properties.  For  example,  magnesium,  zinc,  and  cad- 
mium belong  to  the  same  periodic  group,  but,  by  reason 
of  the  differences  in  the  solubilities  of  their  sulphides, 
the  three  metals  are  placed  in  three  separate  analytical 
groups. 

By  (a)  all  salts  require  two  analyses :  first,  for 
metals ;  and,  second,  for  the  acid  radicals.  By  (b)  and 
(c)  the  metals   are  divided   into  groups  depending  on 


SYSTEMS   OF  ANALYTICAL   EXAMINATION       75 

the  insolubility  of  their  chlorides,  sulphides,  hydroxides, 
and  carbonates. 

There  are  six  groups  of  metals :  — 
Group   I,   whose    chlorides   are   insoluble   in   aqueous 

solution ; 
Group  II,  whose  sulphides  are  insoluble  in  acidified 

(HCl)  solution; 
Group  III,  whose  hydroxides  are  insoluble  in  alkaline 

(NH4OH)  solution ; 
Group  IV,  whose  sulphides  are  insoluble  in  alkaline 

(NH4OH)  solution ; 
Group  V,  whose  carbonates   are  insoluble  in  alkaline 

(NH4OH)  solution ; 
Group  VI,  which  has  no  group  characteristic. 

The  group  reagent  would  be  for :  — 
Group  I,  a  soluble  chloride  (HCl) ; 
Group  II,  a  soluble  acidified  sulphide  (H2S  with  HCl) ; 
Group  III,  a  soluble  alkali  (NH^OH  with  NH^Cl); 
Group  IV,  a  soluble  alkaline  sulphide  [(NH4)2S  with 

NH4OH  and  NH4CI] ; 
Group  V,  a  soluble  alkaline  carbonate  [(NH4)2C03  Avith 

NH4OH  and  NH4CI] ; 
Group  VI,  no  group  reagent. 

By  (b)  and  (d)  the  acids  are  also  divided  into  groups 
depending  on  the  insolubility  of  their  barium  and  silver 
salts,  hence  the  three  groups :  — 
Group  I,  whose  barium  salts  are  insoluble  in  aqueous 

solution ; 
Group   II,  whose  silver  salts  are   insoluble  in  dilute 

acid  solution ; 
Group  III,  which  has  no  group  reagent. 


76  CHEMICAL  ANALYSIS 

By  (b)  and  (<?)  the  individual  metals  and  acid  radicals 
included  in  the  groups  are  separated  and  distinguished. 
For  example,  silver,  lead,-  and  mercury  (mercurous) 
chlorides  are  thrown  down  by  hydrochloric  acid  as 
white  precipitates. 

The  separation  is  accomplished  by  observing  facts 
like  these :  Lead  chloride  is  soluble  in  hot  water ;  sil- 
ver and  mercurous  chlorides  are  not.  Silver  chloride 
is  soluble  in  ammonia ;  white  mercurous  chloride  is  not 
soluble,  but  is  blackened  by  that  reagent.  Thus,  both 
the  differences  of  solubility  and  color  are  used  for  sepa- 
rating and  detecting  metals. 


^  1 


Part  II— Reactions  and  Separations 


CHAPTER  VII 

METALS  OF  GROUP  I :   SILVER,  MERCURY  (MERCUROUS 
SALTS,  Hg'),  AND  LEAD 

Characteristic  :  Insolubility  of  the  chlorides  in  cold  water  or 
dilute  HCl. 

Group  Reagent:  Dilute  hydrochloric  acid. 

REACTIONS 

Silver  (salt  for  study,  silver  nitrate,  AgNOg). 

1.  HCl  precipitates  white  silver  chloride,  AgCl,  dark- 
ening in  the  sunlight;  insoluble  in  HNOg;^  soluble  in 

nh^oh. 

The  soluble  compound  formed  by  adding  NH^OH  to  AgCl 
is  variously  regarded  as  2AgCl-3NH3,  AgCl-2NH3,  and 
AgCl'SNHg,  though  the  preponderance  of  authority  favors 
2AgCl'3NH3.  NH^OH  also  unites  with  many  other  salts, 
especially   those   of   mercury,  copper,    and   cobalt    (which  see). 

1  The  student  is  urged  again  to  be  methodical  and  thoughtful  in  his 
laboratory  exercises,  especially  in  the  execution  of  the  reactions  and 
separations  in  this  part  of  the  book. 

For  convenience  and  economy  of  space,  the  reagents  hereafter  referred 
to  are  generally  expressed  by  molecular  formulas,  —  not  by  names;  but 
the  student  should  not  on  this  account  contract  the  bad  habit  of  calling 
chemical  substances  by  their  formulas.  For  illustration,  "hydrogen  sul- 
phide "  —  not  "  HjS  "  —  should  be  the  spoken  name  for  the  compound. 

77 


78  CHEMICAL  ANALYSIS 

Some  seem  to  be  derivatives  pf  ammonia,  and  others  of  the 
quasi-metal,  ammonium.  (See  Remsen's  Inorganic  Chemistry, 
p.  274.) 

The  theory  that  these  substances  are  chemical  compounds  — 
not  "  molecular  additions  "  —  is  in  harmony  with  their  conduct. 
In  all  of  them  there  has  been  a  shifting  of  the  atoms  of  the 
metallic  salts  and  ammonia  to  form  more  complex  ions,  thus  pro- 
ducing radical  changes  in  their  solubilities  and  chemical  conduct. 

2.  H2S  precipitates  black  silver  sulphide,  Ag2S,  sol- 
uble in  boiling  HNO3. 

3.  NH^OHi  precipitates  brown  silver  oxide,  AggO, 
soluble  in  excess  of  reagent. 

4.  NaOH  gives  similar  results  to  NH^OH,  except 
that  the  precipitate  is  not  soluble  in  excess  of  reagent. 

5.  K2Cr04  precipitates  dark-red  silver  chromate, 
Ag2Cr04,  soluble  in  hot  HNO3  and  in  NH^OH. 

6.  Metallic  copper  deposits  metallic  silver  on  its  surface. 

7.  Reducing  flame  on  charcoal  with  NagCOg,  or,  better, 
with  the  fusion  mixture  of  NagCOg  and  KgCOg,  gives 
a  metallic  silver  bead. 

NagCOg  acts  partly  as  a  flux,  thus  making  possible  the  ioniza- 
tion of  AgNOg  and  NagCOg,  and  partly  as  an  active  chemical 
agent.     The  following  reactions  occur  here  :  — 

2  AgNOg  -h  NagCOg  =  Ag^COg  +  2  NaNOg ; 
Ag^COg  -f  heat  =  Ag^O  +  COg  ; 
2Ag20-f-C=4Ag-f  CO2. 

Mercurous  Mercury,  Hg'  (salt  for  study,  mercurous 
nitrate,  Hg2(N03)2). 

1.  HCl  precipitates  white  mercurous  chloride  or  calo- 
mel, Hg2Cl2,  soluble  in  hot  HNOg ;  insoluble  in 
NH^OH,  forming  a  black  substance. 


METALS    OF  GROUP  I  79 

•  The  black  insoluble  substance  is  aniido-mercurous  cliloride 
("  black  precipitate "),  NHgHg'gCl,  or  possibly  a  mixture  of 
NHgHg^'Cl  ("  white  precipitate  ")  and  metallic  mercury :  — 

Hg2Cl2+  2  NH4OH  =  NH^HgCl  +  Hg  +  NH.Cl  +  2  H^O. 

2.  HgS  precipitates  a  black  mixture  of  mercuric  sul- 
phide^ and  mercury,  HgS  +  Hg,  soluble  in  aqua  regia. 

llm  action  of  aqua  regia  depends  on  the  following  :  — 

3  IICl  +  HNO3  =  CI2  +  NOCl  +  2  H2O  ; 
2  HgS  +  CI2  +  2  NOCl  -  2  IlgCla  +  2  NO  +  2  S. 

3.  NH4OH  produces  a  black  precipitate ^  of  unknown 
composition,  insoluble  in  excess  of  reagent. 

4.  NaOH  produces  a  black  precipitate  of  mercurous 
oxide,  HggO,  insoluble  in  excess  of  reagent. 

5.  SnClj  precipitates  gray  metallic  mercury,  distinct 
when  boiled  with  HCl. 

The  action  of  SnCl.^  on  mercurous  salts  depends  on  the  tend- 
ency of  stannous  salts  to  oxidize  to  the  stannic  condition,  on 
which  account  they  act  as  energetic  reducing  reagents.  Mercurous 
salts  are  reduced  to  metallic  mercury,  and  mercuric  salts  are  first 
reduced  to  mercurous  and  then  to  metallic  mercury :  — 

Hg^Cl^  +  SnClg  =  2  Hg  +  SnCl^ ; 
2HgCl2  +  SnClg  =  HggCla  +  SnCl^  ; 
Hg^Cl^  +  SnCl^  =  2Hg  +  SnCl,. 

6.  KI  precipitates  green  mercurous  iodide,  Hggig, 
soluble  in  acids. 

7.  Metallic  copper  deposits  metallic  mercury,  distinct 
when  polished. 

8.  Heating  in  a  closed  tube,  with  fusion  mixture, 
deposits  globules  of  metallic  mercury  on  the  sides  of 
the  tube. 


80  CHEMICAL  ANALYSIS 

Lead  (salt  for  study,  lead  nitrate,  Pb  (N03)2). 

1.  HCl  precipitates  white  lead  chloride,  PbClg,  solu- 
ble in  boiling  water,  and  in  concentrated  HCl. 

2.  HgS  precipitates  black  lead  sulphide,  PbS,  soluble 
in  hot  HNOg. 

3.  NH4OH  and  NaOH  precipitate  white  lead  hydroxide, 
Pb(0H)2,  soluble  in  excess  of  NaOH,  forming  the 
sodium  salt  NagPbOg.  If  NH^OH  is  used  for  pre- 
cipitation, and  NaOH  for  redissolving  the  precipitate, 
the  latter  should  be  filtered  and  washed  before  add- 
ing the  NaOH.  This  is  necessary  in  order  to  elimi- 
nate ammonium  salts  which  prevent  the  formation  of 
Na2Pb02. 

4.  H2SO4  precipitates  white  lead  sulphate,  PbSO^, 
soluble  in  hot  NaOH  and  in  NH4C2H3O2. 

5.  K2Cr04  precipitates  yellow  lead  chromate,  PbCrO^, 
soluble  in  NaOH,  forming  Na2Pb02. 

6.  KI  precipitates  yellow  lead  iodide.  Pblg,^  soluble 
in  hot  water  and  in  excess  of  KI. 

7.  Metallic  zinc  deposits  metallic  lead. 

8.  Reducing  flame  on  charcoal,  with  fusion  mixture, 
precipitates  metallic  lead. 

PROCESS   OF   SEPARATION 

The  separation  of  the  members  of  this  group  is  based 
upon  the  fact  that  PbClg  is  soluble  in  hot  water, 
and  that  AgCl  is  soluble  in  NH^OH.  The  process 
of  separation  is  as  follows :  — 

Add  cold  dilute  HCl  so  long  as  the  white  precipi- 
tate continues  to  be  formed;  then  add  about  ten  di'ops 


METALS   OF  GROUP  I  81 

in  excess.  Filter  and  wash  ^  with  cold  water  ^  (ice  water 
is  preferable).  Filtrate  (a)  may  contain  members  of 
all  the  subsequent  groups  ;  residue  (a)  may  consist  of 
PbClg,  AgCl,  and  HggGlg.  Pour  boiling  water  on  the 
residue  in  the  filter  ^  and  test  the  filtrate  (b)  with  KgCrO^ 
for  lead,  and,  if  present,  continue  the  washing  with  hot 
water  till  all  traces  of  it  disappear.  Add  warm  NH^OH 
to  the  residue  (b)  on  the  filter  and  test  the  filtrate  (c) 
for  silver^  by  neutralizing  with  HNO3.  Precipitation  of 
AgCl  confirms  the  presence  of  silver.  If  a  black  residue 
(c)  is  left  after  adding  NH^OH,  mercury  is  probably 
present.  Confirm  by  dissolving  in  a  little  aqua  regia, 
evaporating  excess  of  chlorine  and  acids,  and  adding 
SnClg. 

A  certain  amount  of  care  must  be  observed  on  adding  HCl, 
since  dilute  HCl  often  precipitates  certain  members  of  the  next 
group  (BiOCl  and  SbOCl).  This  source  of  error  can  be  removed 
by  adding  a  sufficient  excess  of  HCl,  which  redissolves  the  BiOCl 
andSbOCl. 

On  the  other  hand,  too  great  an  excess  of  HCl  interferes  with 
the  action  of  the  next  group  reagent  (HgS) ;  and,  furthermore,  it 
may  redissolve  the  chlorides  of  Group  I. 


CHAPTER    VIII 

METALS  OF  GROUP  II:  MERCXJRY  (MERCURIC  SALTS, 
Hg"),  LEAD,  BISMUTH,  COPPER,  CADMIUM,  ARSENIC 
(ARSENIOUS,  As'",  AND  ARSENIC,  As^  SALTS),  ANTI- 
MONY,  AND   TIN 

Characteristic:    Insolubility  of  the  sulphides  in  dilute  HCl. 
Group  Reagent:  Hydrogen  sulphide  in  the  presence  of  hydro- 
chloric acid. 

REACTIONS 

Mercuric  Mercury,  Hg"  (salt  for  study,  mercuric  chloride, 
HgCl^). 

1.  HgS  precipitates  black  mercuric  sulphide,  HgS, 
soluble  in  aqua  regia;  insoluble  in  yellow  (NH4)2S^ 
and  in  hot  HNO3. 

When  IlgS  is  first  added,  a  white  precipitate  is  obtained 
which  is  the  salt,  HgClg-  2  HgS.  On  adding  more  HgS,  it  becomes 
orange,  brown,  and  then  black.  If  insufficient  H^S  is  added,  it 
remains  orange  and  is  often  mistaken  for  SbgSg  (which  see). 

2.  NH4OH  precipitates  "  white  precipitate,"  NH2-  HgCl, 
soluble  in  acids. 

3.  NaOH  precipitates  yellow  mercuric  oxide,  HgO, 
soluble  in  acids. 

4.  SnCl2  precipitates  white  Hg2Cl2,  changing  to  gray 
metallic  mercury  by  addition  of  an  excess  of  reagent. 
Boiling  with  a  little  HCl  causes  metallic  globules 
to  form. 

5.  KI  precipitates  scarlet  mercuric  iodide,  Hglg,^  solu- 
ble in  excess  of  reagent,  forming  2KI-Hgl2. 

82 


METALS   OF  GROUP  II  83 

6.  Metallic  copper  deposits  metallic  mercury. 

7.  Heating  in  a  closed  tube  with  fusion  mixture  causes 
separation  of  metallic  mercury. 

Lead.     (See  reactions  for  lead  under  Group  I.) 

Lead,  when  present  in  a  mixture,  is  often  found  in 
Group  II,  because  PbClg  is  somewhat  soluble  in  acidi- 
fied aqueous  solutions,  and  is  brought  over  with  the 
filtrate  from  Group  I. 

Bismuth  (salt  for  study,  bismuth  nitrate,  Bi(N03)3). 

1.  H2S  precipitates  black  bismuth  sulphide,  BigSg, 
soluble  in  boiling  HNO3 ;  insoluble  in  (NH4)2Sx.  The 
precipitation  is  hastened  by  dilution. 

2.  NH4OH  and  NaOH  precipitate  white  basic  bismuth 
hydroxide,  (BiO)OH,  soluble  in  acids.  If  made  alkaline 
again  with  NaOH,  and  SnCl2  is  added,  black  bismuth 
monoxide,  BiO,  will  be  formed. 

3.  Water  added  in  large  excess  precipitates  white 
bismuth  oxynitrate,  (BiO)N03,  soluble  in  acids.  Addi- 
tion of  NaCl  hastens  the  precipitation  by  the  formation 
of  the  less  soluble  (BiO)Cl  by  double  decomposition. 

In  this  experiment,  both  water  and  HNO3  (a  constituent  of 
Bi(N03)3)  endeavor  to  react  with  bismuth,  the  one  to  form 
insoluble  basic  (BiO)N03  and  the  other  to  maintain  the  soluble 
nitrate  61(^03)3.  HNO3  is  far  better  ionized  than  water,  and 
with  a  degree  of  equality  in  the  contest  it  would  maintain  the 
salt  Bi(N03)3  ;  but  the  large  excess  of  water  so  handicaps  the 
acid  that  it  yields  to  mass  action.  In  this  hydrolytic  action 
the  positive  ion,  Bi,  is  changed  to  another  positive  ion,  BiO 
(bismuthyl). 


84  CHEMICAL  ANALYSIS 

4.  Metallic  iron  or  zinc  deposits  metallic  bismuth. 

5.  Reducing  flame  on  charcoal  with  fusion  mixture 
gives  a  brittle  bead  of  metallic  bismuth.  The  border 
of  the  pit  becomes  coated  with  yellow  bismuth  trioxide, 

Copper  (salt  for  study,  copper  sulphate,  CUSO4). 

1.  HgS  precipitates  black  copper  sulphide,  CuS,  solu- 
ble in  HNO3  ^^^  KCN  ;  almost  insoluble  in  (NH4)2Sx ; 
insoluble  in  dilute  H2SO4  and  yellow  NagS^. 

2.  NH4OH  precipitates  a  greenish  basic  cupric  salt, 
probably  CuSO^-  CuO,  soluble  in  excess  of  reagent,  form- 
ing a  blue  cuprammonium  compound.^ 

3.  NaOH  precipitates  blue  copper  hydroxide,  Cu(0H)2, 
darkening  on  boiling  and  giving  SCuO-HgO. 

4.  Metallic  iron  or  zinc  deposits  metallic  copper. 

5.  Reducing  flame  on  charcoal  with  fusion  mixture 
gives  metallic  copper.  The  finely  divided  copper  is 
easily  detected  after  the  reduction  by  triturating  the 
fused  mass  in  a  small  mortar  and  then  washing  off  the 
charcoal  and  dissolved  salts. 

6.  KCN  precipitates  greenish  cupric  cyanide, 
Cu(CN)2,  soluble  in  excess  of  reagent,  forming  the 
salt  CuCN-SKCN  or  K3Cu(CN)4.  This  salt  is  not 
reprecipitated  by  H2S.  This  is  the  reverse  of  Reac- 
tion 1. 

7.  K4Fe(CN)g  precipitates  brown  cupric  ferrocyanide, 
Cu2Fe(CN)g,  insoluble  in  dilute  acids. 


METALS  OF  GROUP  II  85 

CHEMISTRY  OF  CYANOGEN  COMPOUNDS 

As  cyanogen  compounds  are  important  in  many  ana- 
lytical reactions,  it  is  necessary  to  consider  their  behavior 
and  constitution  at  some  length. 

Cyanogen,  CN,  is  much  like  the  halogens,  in  both  the 
constitution  and  behavior  of  its  acids  and  salts.  The 
cyanogen  analogues  of  HCl,  KCl,  HCIO,  KCIO,  etc.,  are, 
respectively,  HCN,  KCN,  HCNO,  KCNO,  etc. 

The  cyanides  of  the  more  basic  metals,  the  alkali  and 
alkali-earth  metals,  are  soluble  and  easily  decomposed  by 
acids,  while  those  of  the  less  basic  heavy  metals  are  less 
soluble  and  more  stable  in  the  presence  of  acids.  Though 
the  cyanides  of  the  strong  basic  metals  are  easily  decom- 
posed by  acids,  they  are  not  readily  dissociated  by  heat 
and  can  be  fused  without  decomposition.  The  reverse  is 
true  of  the  cyanides  of  the  weak  basic  metals.  Many 
insoluble  cyanides  unite  with  soluble  cyanides  to  form 
so-called  "  double  cyanides  "  :  — 

CuCN  +  3 KCN  =  CuCN-SKCN  or  K3Cu(C]Sr)4,  potassium 
cuprous  cyanide ; 

Fe(C]Sr)2  -f  4KCN  =  Fe(CN)2-4KCN  or  K4Fe(CN)6,  potas- 
sium ferrous  cyanide ; 

Fe(C]Sr)3  -f  3 KCN  =  Fe(CN)3-3KCN  or  KgFeCCN)^,  potas- 
sium ferric  cyanide; 

Co(CN)2  -f  4KCN  =  Co(CN)2-4KCN  or  K4Co(CN)6,  potas- 
sium cobaltous  cyanide. 

The  behavior,  with  reagents,  of  the  solutions  of  many  of 
these  combined  salts  indicates  that  they  are  not  mixtures  of 
binary  cyanides  but  salts  of  ternary  acids  whose  acid  radi- 
cals are  cyanides  of  various  metals.  For  convenience  the 
tern:  metallo-cyanide  is  applied  to  this  class  of  salts. 


86  CHEMICAL  ANALYSIS 

Copper  is  usually  precipitated  as  CuS  by  H2S  in  acid 
solution,  but  this  reaction  does  not  occur  if  K3Cu(CN)4  is 
treated  with.  HgS.  This  seems  to  indicate  that  there  are 
no  free  Cu  ions  in  a  solution  of  K3Cu(CN)4.  Potassium 
cyanide  reacting  with  H2S  liberates  hydrocyanic  acid, 
HCN,  but  when  K3Cu(CN)4  is  treated  with  HgS,  HON  is 
not  evolved.  Both  examples  show  that  the  two  constitu- 
ents of  the  salt,  CuCN  and  KCN,  are  so  strongly  united  as 
to  destroy  the  identity  of  each,  and  that  K3Cu(CN)4  is  the 
preferable  formula. 

This  view  is  important  in  analytical  chemistry,  as  it 
interprets  the  behavior  of  several  substances  similar  to 
K3Cu(CN)4. 

The  metallo-cyanides  occurring  in  analytical  chemistry 
may  be  considered  under  three  classes  :  — 

1.  Metallo-cyanides  decomposed  by  hydrogen  sul- 
phide :  — 

K2Cd(CN)4    +  2H2S  =  CdS  +  K^S  +  4HCN ; 
K2Hg(CN)4  -i-  2H2S  =  HgS  +  K2S  +  4HCN-,  etc. 

This  class  is  important  in  the  separation  of  copper  and 
cadmium  in  Group  II  for  metals. 

2.  Metallo-cyanides  decomposed  by  dilute  mineral 
acids  :  — 

K3Cu(C]Sr)4  4-  3HC1  =  CuCN  -f  3KC1  +  3HCN ; 
K2Ni(CN)4  +  2HC1  =  Ni(CN)2  -f  2KC1  +  2HCN ; 
K4Co(CN)6  -f-4HCl  =  Co(CN)2  +  4KCH-4HCN,  etc. 

The  reaction  of  K4Co(CK)6  is  merely  temporary,  and  the 
actual  product  is  K3Co(CN)6,  showing  that  the  cobaltous 
salt  is  oxidized  to  the  corresponding  cobaltic  salt.  The 
nickelous  salt  is  not  oxidized  to  the  nickel ic  salt.  This 
is  further  illustrated  in  the  oxidation  of  K2Ni(CN)4  and 
K4Co(CN)6  by  a  hypobromite:  — 


METALS   OF  GROUP  II  87 

2  K2M(CN)4  +  NaBrO  +  SH^O  =  2Ni(OH)3  +  NaBr 

+  4KCN  +  4HCN; 
2K4Co(CN)6  +  NaBrO  +  H^O  =  2K3Co(CN)6  +  NaBr 

+  2K0H. 

This  class  is  important  in  the  separation  of  nickel  and 
cobalt  in  Group  IV  for  metals. 

3.  Metallo-cyanides  not  decomposed  by  dilute  mineral 
acids. 

The  members  of  this  class  are  stable  chemical  com- 
pounds and  are  not  decomposed  except  by  drastic  treat- 
ment. They  give  a  clearer  insight  into  the  constitution  of 
the  metallo-cyanides.  The  most  important  are  the  three 
reagents,  potassium  ferrocyanide,  K4Fe(CN)6,  potassium 
ferricyanide,  K3re(CN)6,  and  potassium  sulphocyanate, 
KCNS.i  They  react  with  many  compounds  and,  for  the 
most  part,  form  colored  metallo-cyanides:  — 

2CUSO4  4-  K4Fe(CN)6     =  2  K2SO4  +  Cu2Fe(CN)6 ; 
4FeCl3  +  3K4Fe(CN)6   =  12KC1  +  Fe4(Fe(CN)6)8 ; 
3FeS04  H-  2K3Fe(CN)6  =  3K2SO4  -f-  Fe3(Fe(CN)6)2; 
FeCl3  +  3  KCNS  =  3  KCl  -f  Fe(CNS)3. 

This  class  is  important  in  the  detection  of  copper  and 
ferrous  and  ferric  iron. 

Cadmium  (salt  for  study,  cadmium  nitrate,  Cd(N0)3)2. 

1.  HgS  precipitates  yellow  cadmium  sulphide,  CdS^ 
very  soluble  in  boiling  acids;  insoluble  in  cold  dilute 
acids,  (NH4)2S^,and  KCN. 

2.  NH4OH  precipitates  white  cadmium  hydroxide, 
Cd(0H)2,  soluble  in  excess  of  reagent. 

3.  NaOH  acts  like  NH^OH,  but  does  not  redissolve 
Cd(0H)2  in  excess  of  reagent. 


88  CHEMICAL  ANALYSIS 

4.  Metallic  zinc  deposits  metallic  cadmium. 

5.  Reducing  flame  on  charcoal  with  fusion  mixture 
gives  a  brown  incrustation  of  cadmium  oxide,  CdO. 
The  salt  is  first  reduced  to  metallic  cadmium,  which, 
owing  to  its  ready  volatility,  rises  to  a  stratum  of  oxy- 
gen and  is  oxidized. 


Arsenious  Arsenic,  As'"  (salt  for  study,  sodium  arsenite, 

Na3As03). 

1.  H2S  precipitates  from  acidified  solutions,  yellow 
arsenious  sulphide,  AsgSg,  soluble  in  (NH4)2S^,(NH4)2C03, 
and  aqua  regia  ;  insoluble  in  concentrated  boiling  HCl. 

2.  AgNOg  precipitates  from  very  weak  ammoniated 
solutions,  yellow  silver  arsenite,  AggAsOg,  soluble  in 
excess  of  NH4OH  and  HNO3. 

3.  CUSO4  precipitates  from  ammoniated  solutions, 
Scheele's  green,  HCuAsOg,  soluble  in  excess  of 
NH4OH. 

4.  Heating  in  a  closed  tube  with  NagCOg  and  KCN 
deposits  a  mirror  of  metallic  arsenic  on  the  cold  part  of 
the  tube. 

5.  Nascent  hydrogen  reduces  arsenic  compounds  to 
arsine,  AsHg.      If   the   hydrogen   and   the  AsHg  are 

'passed  through  a  jet  tip  and  kindled,  AsHg  will  first 
be  oxidized  to  metallic  arsenic  and  then,  on  emerging 
from  the  flame,  will  be  further  oxidized  to  AsgOg.  If 
AsHg  is  passed  into  a  solution  of  AgNOg,  metalHc 
silver  and  arsenious  acid,  Hg  AsOg,  will  be  produced :  — 

6  AgNOg  +  AsHg  +  3  H2O  =  6  Ag  +  6  IINOg  +  Hg  AsOg. 


METALS   OF  GROUP  II  89 

These  reactions  constitute  the  principles  of  — 

Marsh's  Test^for  Arsenic.  —  Generate  hydrogen  in  a 
small  flask  with  pieces  of  pure  zinc  or  magnesium  and 
dilute  H2SO4  or  HCl.  The  flask  should  have  a  funnel 
tube,  and  the  horizontal  delivery  tube  should  have  a 
wide,  hard  glass  tube  about  2  cm.  in  diameter  and  3  dm. 
long,  drawn  to  a  jet  point,  and  with  two  constrictions 
near  the  middle.  When  the  zinc  is  put  into  the  flask 
and  all  connections  made  tight,  the  dilute  acid  is 
poured  through  the  funnel  tube.  After  two  or  three 
minutes,  test  the  hydrogen  by  collecting  by  displace- 
ment of  water  in  a  test-tube.  If  it  kindles  quietly 
with  a  blue  flame,  the  jet  may  now  be  lighted.  To 
insure  safety  a  cloth  should  be  thrown  over  the  appa- 
ratus. Test  the  reagents  and  apparatus  for  arsenic  by 
holding  a  cold  porcelain  dish  at  the  top  of  the  flame. 
If  no  black  spots  appear,  no  arsenic  is  present. 

A  concentrated  HCl  solution  of  the  arsenic  com- 
pound is  poured  through  the  funnel.  Arsenic  can  be 
detected  in  four  ways  :  — 

{a)  The  hydrogen  flame  assumes  a  pale  violet  color, 
due  to  the  oxidation  of  AsHg  to  white  AS2O3. 

(h)  Black  spots  of  metallic  arsenic  are  deposited  on 
a  cold  porcelain  plate.  The  plate  should  not  be  heated 
too  much,  as  the  arsenic  would  then  be  sublimed.  The 
black  spots  can  be  removed  by  a  solution  of  sodium 
hypochlorite  :  — 

lONaClO  +  4  As  +  6H2O  =  lONaCl  -f-  4H3ASO4. 

(c)  If  the  hard  glass  tube  is  heated  at  one  of  the  con- 
tractions, the  heat  will  decompose  the  passing  AsHg 


90  CHEMICAL   ANALYSIS 

and  will  cause  metallic  arsenic  to  be  deposited  on  the 
cold  part  of  the  tube.  When  the  evolution  of  hydro- 
gen is  completed,  detach  the  tube  with  the  arsenic 
mirror  from  the  generator  and  pass  HgS  through  the 
slightly  warmed  tube.  The  metallic  mirror  will  change 
to  yellow  AS2S3. 

(d)  Hofmanns  Modification.  —  Instead  of  kindling 
AsHg,  pass  it  through  a  U-tube  containing  glass 
splinters  or  beads  moistened  with  a  solution  of  lead 
acetate,  in  order  to  counteract  any  passing  HCl  or  H2S, 
and  then  pass  the  gas  into  a  test-tube  containing  a 
solution  of  AgNOg.  When  the  precipitation  is  com- 
pleted, remove  the  test-tube  and  add  very  cautiously 
a  thin  stratum  of  very  dilute  NH^OH.  The  junction 
of  the  two  liquids  should  show  a  faint  yellow  ring  of 
AggAsOg. 

Arsenic  Arsenic,  As^  (salt  for  study,  sodium  arsenate, 
NagAsO^). 

1.  H2S  precipitates,  from  strong  HCl  solutions  (2 
parts  concentrated  HCl  to  1  part  water),  yellow  arsenic 
sulphide,  AsgS^.  HgS  will  precipitate  AsgSg  from 
dilute  solutions  of  HCl  if  they  are  warmed  to  about 
70°  and  the  gas  is  passed  through  for  some  time. 
AsgSg  behaves  much  like  A 8383  towards  (NH4)2S^,  aqua 
regia,  and  (NH4)2C03. 

ASgSg  is  a  colloidal  precipitate  and  requires  for  its  precipita- 
tion both  heat  and  strong  acid  solution.  (See  Washing,  p.  29.) 
The  semi-solid  often  colors  the  solution  yellow  without  precipi- 
tation ;  in  which  event  it  is  necessary  to  add  more  HCl  and  in- 
crease the  heat.     Furthermore,  IlgS  with  cold  dilute  HCl  does 


METALS   OF  GROUP  II  91 

not  precipitate  AsgSg  very  well,  as  the  arsenic  is  reduced  to  the 
-ous  condition :  Na^AsO^  +  IlgS  =  NagAsOg  +  U^O  +  S. 

2.  AgNOg  precipitates  from  a  very  slightly  alkaline 
solution,  brown  silver  arsenate,  AggAsO^,  soluble  in 
excess  of  NH^OH  and  HNO3. 

3.  MgSO^  in  presence  of  NH^Cl  and  NH^OH  precipi- 
tates ammonium  magnesium  arsenate,  NH^MgAsO^.^ 

The  purpose  of  adding  NH^Cl  is  to  avoid  the  formation  of 
the  insoluble  compound  Mg(0H)2;  and  the  purpose  of  adding  an 
excess  of  NH^OH  is  to  insure  the  insolubility  of  the  double  salt 
NH.MgAsO^. 

4.  (NH4)2Mo04  in  HNO3  solution  precipitates  yel- 
low ammonium  arseno-molybdate,  (Mo03)j2(NH^)3As04. 
The  precipitation  does  not  occur  in  the  cold,  like 
that  from  phosphoric  acid  (which  see). 

(NH4)2Mo04  in  IINOg  solution  changes  to  the  more  com- 
plex salt  (Nri^)2Mo^Oj3,  wliicli,  with  IlgAsO^,  forms  insoluble 
(Mo03)i2'(NIl4)3As04.  The  reagent  should  be  added  in  large 
excess  (about  4  :  1),  as  the  precipitate  is  soluble  in  IlgAsO^. 

5.  Similar  to  4,  under  arsenious  compounds. 

6.  Similar  to  5,  under  arsenious  compounds. 

Antimony  (salt  for  study,  antimonious  chloride,  SbCl3). 

1.  H2S  precipitates  orange  antimonious  sulphide, 
SbgSg,  soluble  in  (NH4)2Sx,  hot  concentrated  HCl, 
and  aqua  regia;   insoluble  in  (NH4)2C03. 

2.  NH4OH  and  NaOH  precipitate  white  antimonious 
hydroxide,  Sb(0H)3,  soluble  in  excess  of  NaOH. 

3.  AgNOg  precipitates  from  alkaline  solutions  of  anti- 
monious salts  a  black  mixture  of  Ag^O  and  metallic 


92  CHEMICAL  ANALYSIS 

silver.  From  antimonic  salts,  AgNOg  precipitates 
white  silver  antimonate,  Ag3Sb04,  soluble  in  NH^OH. 
These  reactions  are  useful  for  detecting  the  -ous  and 
the  -de  conditions  of  antimony. 

4.  Heating  on  charcoal  with  fusion  mixture  gives  a 
metallic   antimony   bead   with   white   incrustation    of 

5.  Water  in  large  excess  hydrolyzes  soluble  antimo- 
nious  salts  and  forms  white  insoluble  basic  salts :  — 

SbClg  +  HgO  =  SbOCl  +  2  HCl. 

6.  Nascent  hydrogen  reduces  antimony  compounds  to 
stibine,  SbHg,  and  deposits  metallic  antimony  on  kin- 
dling, like  arsenic.  If  SbHg  is  passed  into  a  solution 
of  AgNOg,  black  silver  antimonide,  SbAgg,  will  be  pre- 
cipitated :  — 

3  AgNOg  +  SbHg  =  SbAgg  +  3  HNOg. 

On  account  of  the  likeness  between  the  reactions  of 
ai*senic  and  antimony  with  nascent  hydrogen,  it  is  well 
to  compare  the  behaviors  of  the  two  metals  in  Marsh's 
test. 

{a)  The   hydrogen   flame   is   not    tinted    violet    by 

SbHg. 

(b)  The  black  antimony  spots  on  cold  porcelain  are 
blacker  and  less  lustrous  than  those  of  arsenic.  Anti- 
mony spots  are  not  removed  by  hypochlorite  solutions. 

(c)  H2S  passed  through  the  hard  glass  tube  contain- 
ing antimony  stain  produces  orange  SbgSg. 

(d)  In  Hofmann's  test,  SbHg  passed  into  AgNOg  solu- 
tion precipitates  black  SbAgg.  When  the  precipitation 
is  completed,  decant  the  liquid  and  wash  the  residue  by 


METALS   OF  GROUP  IL  93 

decantation.  Dissolve  the  antimony  by  boiling  with 
a  strong  solution  of  tartaric  acid  to  which  a  few  drops 
of  HNO3  have  been  added.  Filter,  acidify  with  HCl, 
and  pass  H2S  through  the  warmed  solution.  An 
orange  precipitate  indicates  Sb2S3. 

Stannous  Tin,  Sn"  (salt  for  study,  stannous  chloride, 
SnClg). 

1.  H2S  precipitates  brown  stannous  sulphide,  SnS, 
soluble  in  (NH4)2Sx  and  hot  concentrated  HCl ;  insolu- 
ble in  dilute  cold  HCl  and  (NH4)2C03. 

2.  NH4OH  and  NaOH  precipitate  white  stannous 
hydroxide,  Sn(0H)2,  soluble  in  excess  of  NaOH, 
forming   sodium   stannite,    NagSnOg. 

3.  HgCl2  freely  added  precipitates  white  HggClg. 

4.  Metallic  zinc  deposits  metallic  tin. 

5.  Heating  on  charcoal  with  fusion  mixture  gives 
metallic  tin  and  a  white  incrustation  of  SnOg. 

Stannic  Tin,  Sn'^  (salt  for  analysis,  stannic  chloride, 
SnCl4). 

1.  H2S  precipitates  yellow  stannic  sulphide,  SnSg, 
soluble  in  (NH4)2Sx  and  hot  concentrated  HCl ;  insolu- 
ble in  dilute  cold  HCl  and  (NH4)2C03. 

2.  NH4OH  and  NaOH  precipitate  white  stannic  hydrox- 
ide, SnO(OH)2,  soluble  in  excess  NaOH. 

3.  Similar  to  4,  under  stannous  salts. 

4.  Similar  to  5,  under  stannous  salts. 


94  •    CHEMICAL  ANALYSIS 

PROCESS  OF   SEPARATION 

The  group  is  divided  into  two  sub-groups.  Sub- 
group A',  arsenic,  antimony,  and  tin,  —  metals  whose 
sulphides  are  soluble  in  (NH4)2Sx.  Sub-group  B :  mer- 
cury, lead,  bismuth,  copper,  and  cadmium,  —  metals 
whose  sulphides  are  not  soluble  in  (NH^gS^. 

The  separation  of  the  members  of  Sub-group  A 
depends  upon  the  following  properties  :  — 

(a)  Insolubility  of  arsenic  sulphide  in  boiling  HCl, 
or  the  solubility  of  arsenic  sulphide  in  (NH4)2C03 
solution. 

(b)  Insolubility  of  metallic  antimony  in  dilute  HCl. 

(c)  Hofmann's  separation  of  arsenic,  antimony,  and 
tin. 

The  separation  of  Sub-group  B  depends  upon  the 
following  properties  :  — 

(a)  Insolubility  of  HgS  in  boiling  dilute  HNO3  ; 

(b)  "  "•  PbSO^  in  acidified  solution ; 

(c)  "  "  (BiO)OH  in  NH4OH  solution; 
(6^)           "  "  CuS  in  dilute  HgSO^,  or 

"  "  CdS  in  KCN  solution. 

Into  the  solution  1  acidified  with  HCl^  and  warmed  to 
about  70°,  pass  a  constant  stream  of  HgS  for  about 
fifteen  minutes ;  then  cool  the  diluted  solution,  and 
before  filtering  pass  HjS  again  till  the  precipitation 
is  completed.  The  precipitation  of  AsgSg  and  AsgS^ 
requires  heat  and  a  strongly  acid  (HCl)  solution,  both 
of  which  tend  to  dissolve  the  other  sulphides.  Hence 
it  is  necessary  to  cool  and  dilute  the  solution  in  order 
to  precipitate  all  the  sulphides.^    Filter.    The  filtrate  (a)  ^ 


METALS   OF  GROUP  II  95 

may  contain  members  of  subsequent  groups.  The  resi- 
due (a)  may  consist  of  the  sulphides  of  all  the  members 
of  the  group.  Thoroughly  wash  the  residue  with  the 
aid  of  the  suction  pump,  remove  it  from  the  paper 
and  digest  it  with  (NH4)2S^i  (or  with  NagS^,^  if  the  pres- 
ence of  copper  is  suspected)  for  about  ten  minutes. 
Filter  and  wash  the  residue  with  water  containing 
some  (NH4)2Sx,  rejecting  the  washings.  The  filtrate 
(b)  may  contain  the  sulpho-salts  of  Sub-group  A.  The 
residue  (b)  may  consist  of  the  sulphides  of  Sub-group  B. 

Separation  of  Sub-group  A.  —  To  the  filtrate  (b)  add 
dilute  HCl  till  acid.  Filter  and  reject  the  filtrate  and 
washings.  A  yellow  precipitate,  residue  (c),  indicates 
the  presence  of  members  of  the  sub-group,  or  it  may  be 
sulphur  or  a  mixture  of  the  sulphides  and  sulphur. 
Sulphur  can  be  dissolved  by  shaking  the  precipitate 
with  benzol  or  petroleum  ether.  If  all  the  precipitate 
dissolves,  it  is  only  separated  sulphur.  Two  methods 
can  be  used  for  the  further  seijaration  of  Sub-group  A. 

Method  1.  —  Treat  residue  ((?),  possibly  consisting  of 
the  reprecipitated  sulphides,  with  warm  concentrated 
HCl  (concentrated  HCl  and  a  little  water).  Filter. 
The  filtrate  (aJ)  may  contain  antimony  and  tin  chlorides. 
The  residue  {a')  may  bfe  sulphur  and  AsgSg.  Divide 
the  residue  into  two  parts.  First  part :  Dry  the  yellow 
mass  and  heat  in  a  closed  tube  with  NagCOg  and  KCN. 
A  mirror  on  the  sides  of  the  tube  indicates  arsenic. 
Second  part:  Fuse  with  Na2C03  on  a  platinum  foil; 
dissolve  in  HNOg  and  boil  to  expel  carbonic  acid. 
Add  excess  of  (NH4)2MoO/ solution  in  HNOg  and  boil. 
A  yellow  precipitate  confirms  the  presence  of  arsenic. 


96  CHEMICAL  ANALYSIS 

Boil  filtrate  (a/)  to  expel  HgS.  Transfer  to  a  small 
dish,  and  immerse  in  the  solution  a  galvanic  couple  of 
strips  of  zinc^  and  platinum.  After  the  evolution  of 
gas,  black  antimony  coats  the  platinum  and  gray  tin 
loosely  adheres  to  the  zinc.  Wash  the  platinum  and 
boil  with  tartaric  acid  and  a  drop  of  fuming  HNO3. 
HgS  producing  an  orange  precipitate  confirms  anti- 
mony. With  the  fingers  rub  off  the  loose  tin  from 
the  zinc 2  into  a  dish.  Again  introduce  the  platinum 
and  some  HCl,  and  boil  till  all  the  loose  tin  dissolves. 
(Some  insoluble  particles  of  carbon  may  remain.) 
Filter,  if  necessary,  and  pour  into  a  test-tube  rinsed 
with  HgClg.  A  white  precipitate  turning  gray  on 
boiling  confirms  tin. 


CHEMISTRY   OF   SULPIIO-COMPOUNDS 

The  solution  of  the  sulphides  of  arsenic,  antimony,  and  tin  by 
(NH4)2Sx  is  due  to  the  formation  of  the  sulpho-salts  of  these  ele- 
ments. They  occur  in  the  periodic  system  about  midway  between 
the  acid-producing  and  base-producing  elements.  Hence  their 
oxides  are  generally  basic  with  reference  to  strong  acid  oxides, 
and  acid  with  reference  to  strong  basic  oxides.  In  the  presence 
of  strong  bases,  the  oxides  AsgOg,  AsgOg,  SbgOg,  SbgOg,  SnO,  and 
SnOg  form  classes  of  -ous  and  -ic  salts :  — 

AsgOg  +  6  NaOH  =  3  R^O  +  2  NagAsOg,  sodium  arsenite ; 

AsgOg  -f-  2NaOn  =     II2O  +  2]SraAs02,        "  metarsenite  ; 

AS2O5  -f  6  NaOH  =  3  HgO  +  2  NagAsO^,       "  arsenate ; 

SbgOg  -f  6  NaOH  =  3  HgO  -t-  2  NagSbOg,      «  antimonite ; 

SbgOg  4-  2]SraOH  =    H2O  4-  2  NaSbOg,        "  metantimonite ; 

Sb205  -j-  6  NaOH  =  3  H2O  +  2  NagSbO^,       "  antimonate ; 

2  SnO  -f  2NaOII  =    HgO  +  Na2Sn203(?),  "  oxystannite; 

SnOg    +  2NaOII  =    HgO  +  NagSnO.,         "  stanuate. 


METALS   OF  GROUP  II  97 

The  corresponding  soluble  sulpho-  or  thio-salts  are  made  in 
two  ways :  — 

First :  Reactions  of  H2S  with  the  oxy-salts  of  the  elements :  — 

Nag AsOg  +  3  HgS  =  3  HgO  +  Nag AsSg,  sodium  sulpharsenite ; 
NaAsOg     -f  2H2S  =  2H2O  +  NaAsSg,       "       metasulpharsenite; 
NagAsO^  +  4  H2S  =  4H2O  +  NagAsS^,      "        sulpharsenate ;  etc. 

When  the  sulpho-salts  are  treated  with  HCl  the  hypothetical 
sulpho-acids  are  formed,  but  are  immediately  broken  down  to  the 
insoluble  sulphides.  This  is  one  reason  that  the  presence  of  HCl 
is  necessary  for  the  precipitation  of  the  sulphides  by  HgS. 

Second:  Reactions  of  alkali  sulphides  with  the  sulphides  of 
the  elements :  — 

AS2S3  +  3(NH4)2S  =  2(NH,)3AsSg; 

AS2S3  +    (NHJ2S  =  2  NH4ASS2 ; 

AS2S5  +  SCNHJgS  =  2(NHj3AsS,; 

Sb2S3  +  3(NHJ2S  =  2(NHj3SbS3; 

Sb2S3  4-    (NH4)2S  =  2  NH4SbS2 ; 

Sb2S5  +  3(NHJ2S  =  2(NH4)3SbS4; 

SnS     +    (NH4)2S  =  no  sulpho-salt  formed ; 

SnS2    +    (NHJ2S  =  (NHj2SnS3. 

All  of  the  higher  sulphides  readily  unite  with  (NH4)2S  to  form 
-ate  salts,  but  the  lower  sulphides  unite  with  (NH4)2S  with  vary- 
ing difficulty. 

AsgSg  reacts  with  (NH4)2S  very  readily ; 
SbgSg      "         "  "      .  with  difficulty ; 

SnS         "         "  "         not  at  all. 

Hence,  in  order  to  change  all  the  insoluble  sulphides  to  solu- 
ble sulpho-salts,  it  is  necessary  to  add  ammonium  polysulphide, 
(NH^)2Sx,  so  as  to  supply  sufficient  sulphur  to  convert  all  of  the 
sulphides  to  the  -ate  salts  : — 

AS2S3  +  3(NH,)2S,  =  2(NH4)3AsS,  +  (3  X-  5)S  ; 
AS2S5  +  3(NHj2Sx  =  2(NH,)gAs  S^  +  3(X  -  1)S; 
Sb2S3  +  3(NH,)2Sx  =  2(NHj3SbS4  +  (3X  -  5)S  ; 
•  SnS  +  (NH4)2Sx=  (NH4)2SnS3  +  (X  -  2)S ; 
SnSa  +    (NHj2Sx=    (NHJ^SnSg   +    (X-l)S. 


98     .  CHEMICAL  ANALYSIS 

As  stated  above,  these  soluble  sulpho-salts  are  readily  decom- 
posed by  HCl  and  the  sulphides  reclaimed  :  — 

2(NHj3AsS4  +  6  HCl  =  As^S^  +  G  NH^Cl  +  3  H^S  ; 
2(NIl4)3SbS4  +  G  HCl  =  SbgSg  +  G  NH^Cl  +  3  HgS,  etc. 

Method  ^,  HofmanrCs  Separation.  —  Dissolve  the 
mixed  sulphides  with  HCl  and  a  crystal  of  KClOg, 
and  evaporate  the  chlorine  and  excess  of  acids.  Ar- 
range the  apparatus  for  Hofmann's  modification  of 
Marsh's  Test  and  add  the  dissolved  sulphides  to  the 
contents  of  the  generating  flask.  When  the  evolution 
of  the  gas  has  ceased,  filter  the  contents  of  the  tes1> 
tube.  The  black  residue  may  consist  of  metallic  silver 
and  SbAgg.  Test  for  SbAgg  by  dissolving  in  tartaric 
acid,  etc.  The  filtrate  may  contain  the  hypothetical 
acid  HgAsOg  with  excess  of  AgNOg,  which  can 
be  tested  by  producing  yellow  AggAsOg  with  dilute 
NH4OH.  Filter  the  contents  of  the  generating  flask, 
remove  the  undissolved  zinc,  and  test  the  residue  for 
tin  by  dissolving ,  in  a  small  amount  of  HCl  and  then 
adding  HgCl2. 

The  reactions  involved  in  the  separation  of  arsenic,  antimony, 
and  tin  by  Hofmann's  method  are  for  :  — 

(a)  arsenic,  —  AsHg  +  3  HgO  +  6  AgNOg  =  3  Agg  +  G  HNOg 

+  H3ASO3; 
(h)  antimony,  — SbHg  +  3  AgNOg  =  SbAgg  +  3  HNO3; 
(c)  tin,  — SnCl^  +  4  H  =  Sn  +  4HC1. 

Separation  of  Sub-group  B.  —  Transfer  the  residue  (6), 
supposed  to  consist  of  the  members  of  the  sub-group, 
to  an  evaporating  dish,  and  boil  with  HNOg^  (diluted 
1 : 2)  till    the    chemical    action    ceases.      Filter.      The 


METALS   OF  GROUP  II  .  99 

residue  {a')'^  may  be  black  HgS  or  a  mixture  of  HgS 
and  white  Hg(N03)2*  2  HgS.  Dissolve  in  aqua  regia, 
boil  off  the  chlorine  and  excess  of  acids,  and  test 
with  SnCl2.  The  filtrate  (a)  may  contain  Pb(N03)2, 
Bi(N03)3,  Cu{N03)2,  and  Cd(N03)2.  Concentrate  this 
filtrate  until  most  of  the  HNO3  is  driven  off,  add 
dilute  H2S04,2  warm  gently,  and  allow  to  stand  for 
some  time.  A  white  precipitate  is  PbSO^.^  If  lead  is 
present,  add  an  excess  of  dilute  H2SO4  and  evaporate 
till  all  the  HNO3  is  expelled.  Dilute  with  water,  place 
aside  to  enable  the  precipitate  to  settle,  and  filter  off 
the  insoluble  residue  (b').  Test  it  by  boiling  with 
NH^C2H302  and  adding  K2Cr04.  A  yellow  precipi- 
tate confirms  PbCrO^.  If  lead  is  present,  use  the 
filtrate  from  PbSO^;  if  lead  is  absent,  boil  off  excess 
of  HNO3  from  filtrate  (a').  Add  NH4OH  till  alkaline. 
A  blue  fluid  confirms  the  presence  of  copper  and  a 
white  flocculent  precipitate,  bismuth.  Filter.*  Dissolve 
the  residue  (c')  in  a  few  drops  of  HCl  and  test 
for  BiOCP  by  adding  an  excess  of  water.  If  the 
filtrate  (c')  is  blue,  add  a  dilute  solution  of  KCN  care- 
fully till  the  color  disappears.  Often  copper  is  present 
in  small  quantities,  and  the  blue^  color  is  not  distinct. 
In  this  event  evaporate  a  small  quantity  of  the  filtrate 
almost  to  dryness,  acidify  with  very  dilute  HCl,  and 
add  K4Fe(CN)g.  A  brown  precipitate  or  coloration 
indicates  Cu2Fe(CN)g.  If  copper  is  present,  add  very 
little  dilute  KCN  solution.  If  copper  is  absent,  adding 
KCN  is  not  necessary.  Pass  HgS.  A  yellow*^  precipi- 
tate confirms  CdS. 


CHAPTER   IX 

METALS  OF  GROUP  III :  ALUMINUM,  CHROMIUM,  AND  IRON 

Characteristic  :  Insolubility  of  the  hydroxides  in  alkaline 
(NH4OH)  solution  in  the  presence  of  ammonium  chloride. 

Group  Reagent  :  Ammonium  hydroxide  with  ammonium 
chloride. 

REACTIONS 
Aluminum  (salt  for  study,  aluminum  sulphate,  Al2(S04)3). 

1.  NH4OH  precipitates  gelatinous  aluminum  hydrox- 
ide, A1(0H)3,  soluble  somewhat  in  excess  of  reagent  in 
the  cold,  but  wholly  insoluble  if  NH4CI  is  present  or  if 
the  solution  is  boiled. 

2.  NaOH  acts  like  NH4OH,  except  that  A1(0H)3  is 
completely  dissolved  in  excess  of  reagent,  forming 
sodium  aluminate,  NagAlOg.  This  in  turn  is  recon- 
verted into  insoluble  A1(0H)3  if  the  solution  is  boiled 
with  NH4CI. 

3.  BaCOg,  suspended  in  water,  precipitates  aluminum 
completely  in  the  cold  as  A1(0H)3  mixed  with  a  basic 
salt,  probably  A1(0H)C03. 

4.  (NH4)2S  precipitates  Al(OH)g  with  evolution  of  HgS. 

5.  HNa2P04  precipitates  white  aluminum  phosphate, 
AlPO^-HgO,  soluble  in  alkalies  in  the  absence  of 
NH4CI  and  in  HCl  and  HNO3 ;  insoluble  in  HC2H3O2. 

6.  Spectrum  (see  Special  Method  for  Aluminum, 
p.  62). 

100 


METALS   OF   GROUP  III  101 

INFLUENCE    OF  AMMONIUM  SALTS  ^ 

The  reactions  between  many  salts  and  alkalies  in  the  presence  of 
ammonium  compounds  demand  the  following  further  explanation. 
•  Soluble  alkalies  react  with  salts  of  many  metals  to  produce 
hydroxides  of  various  solubilities.  Some  of  the  more  insoluble 
of  these  behave  as  weak  acids  in  the  presence  of  strong  bases  and 
combine  with  them  to  form  new  classes  of  soluble  salts. 
The  following  is  a  type  of  this  class  of  reactions  :  — 

ZnCla  +  2  NaOH  =  Zn(0H)2  +  2  NaCl ; 
Zn(0H)2  +  2  NaOH  =  Na2Zn02  (sodium  zincate)  +  2H2O. 

Ammonia  behaves  much  like  the  other  soluble  bases,  provided 
certain  precautions  are  observed.  The  following  reactions  can 
occur  :  — 

ZnCl2  +  2  NH4OH  =  Zn(0H)2  +  2  NH^Cl ; 
Zn(0H)2  +  2NH40H  =  (NH4)2Zn02(ammonium  zincate)  +  2  H2O. 

The  last  reaction  is  interrupted  in  two  ways :  — 

First:  NH3  splits  off  easily,  thus  allowing  Zn(0H)2  to  be 
reclaimed.  Second:  the  by-product,  NH^Cl,  strongly  influences 
the  reaction  and  redissolves  the  precipitated  Zn(H0)2. 

In  order  to  obtain  (Isril4)2Zn02  it  is  necessary  to  add  NH^OII 
to  a  cold  solution  of  ZnCl2,  to  avoid  decomposition  of  (NH4)2Zn02 
into  NHg  and  Zn(0H)2,  and  to  filter  oif  the  solution  of  NH^Cl. 
When  an  excess  of  NH^Cl  or  other  ammonium  salt  is  added,  all 
of  these  conditions  of  solubility  are  modified.  This  applies  to 
the  reactions  with  the  hydroxides  of  the  alkali  metals  as  well 
as  ammonia. 

In  the  case  of  the  hydroxides  of  the  bivalent  metals,  Fe(0H)2, 
Zn(0H)2,  Mn(0H)2,  Co(OH)2,  Ni(0H)2,  Ba(0H)2,  Sr(0H)2, 
Ca(0H)2,  and  Mg(0H)2,  the  tendency  of  the  alkalies  to  redissolve 
them  is  greatly  augmented  by  NH^Cl.  Hence  these  hydroxides 
are  not  precipitated  in  the  presence  of  ammonium  salts. 

In  the  case  of  the  hydroxides  of  the  trivalent  metals,  A1(0H)3, 
Cr(0II)3,  and  Fe(0H)3,  the  tendency  of  the  alkalies  to  redissolve 
them  is  counteracted  by  NH^Cl. 


102  CHEMICAL  ANALYSIS 

Hence  these  hydroxides  are  completely  precipitated  in  the 
presence  of  ammonium  salts. 

Two  theories  are  now  sanctioned  by  good  authorities  for  the 
interpretation  of  the  influence  of  ammonium  salts  in  the  reactions 
just  mentioned. 

1.    Double  Salts  Theory 

(a)  Salts  of  Bivalent  Metals.  —  The  following  pairs  of  reac- 
tions will  explain  this  theory  with  reference  to  the  salts  of  biva- 
lent metals :  — 

ZnClg  +  2NaOH  (or  NH^OH)  =  Zn(0H)2  +  2NaCl, 
Zn(0II)2  +  4NH4C1=  ZnClVSNH^Cl  (a  soluble  double  salt) 

+  2NH4OII; 
MgClg  +  2  NaOH  =  Mg(0II)2  +  2NaCl,  . 
Mg(OH)2  +  3NH4CI  =  MgCl2-  NH4CI  +  2NH4OH. 

(b)  Salts  of  Trivalent  Metals.  —  The  application  of  equations 
analogous  to  those  of  the  preceding  paragraph  would  lead  us  to 
expect  the  following  reactions  to  occur  with  hydroxides  of  triva- 
lent metals :  — 

t  (NH4)3A103  +  3  HCl  +  NH.Cl, 
Al(OH)3  +  4NIl4Cl-:-]  or 

(        AICI3  •  NH4CI  +  3  NH4OH. 

But  the  hypothetical,  soluble  bodies,  (NH4)3A103  and 
AlClg-NH^Cl,  are  not  known  to  exist,  and  their  non-existence 
is  taken  to  explain  the  failure  of  the  hj'^droxides  of  trivalent 
metals  to  dissolve  in  the  solutions  of  ammonium  salts. 

2.    Ionic  Theory 

Another  explanation  of  the  part  played  by  ammonium  salts  is 
based  upon  the  simple  ionic  principle  that  the  addition  of  an  ion 
in  common  with  one  in  the  solute  decreases  the  dissociation  of 
the  latter. 

(a)  Salts  of  Bivalent  Metals.  —  When  NaOH  or  NII^OH  is 
added  in  excess  to  the  solution  of  a  bivalent  metal,  a  precipitate 


METALS   OF  GROUP  III  103 

is  formed  which  readily  dissolves  on  the  addition  of  NH^Cl.  The 
explanation  is  that  the  addition  of  the  common  ion,  NH^,  sup- 
presses the  negative  ion,  OH,  thus  driving  the  dissociated  ions 
into  undissociated  and  inactive  molecules  of  NH^OH.  As  the 
hydroxides  of  the  bivalent  metals  in  question  are  usually  moder- 
ately well  dissociated,  NH^Cl  would  not  only  suppress  the  free 
OH  ions  of  any  excess  of  NH^OH,  but  also  those  of  the  hydrox- 
ides themselves.  Of  course  it  must  be  understood  that  if  NaOH 
is  used,  this  reaction  first  occurs :  — 

NaOH  +  NH.Cl  =  NaCl  +  NH^OH. 

For  application  of  the  principle  the  important  case  of  magne- 
sium salts  with  NH^OH  and  NH^Cl  is  considered :  — 
Substituting  in  the  equation  a'b  =  ck,^  the  equation 

NH,  X  OH  =  NH.OH  x  k,  or  ^^^'  ^^,?/^  =  k,  is  obtained. 

NH^,  OH,  NH^OH,  and  k  are,  respectively,  the  positive  and  nega- 
tive ions,  the  undissociated  molecules,  and  the  ionization  constant 
for  NH^OH.  Now  when  NH^Cl  is  introduced  the  number  of 
NH^  ions  is  greatly  increased,  and  the  result  is  to  suppress  the 
OH  ions.  By  letting  x  =  number  of  NH^  ions  added,  and  y  = 
those  of  the  undissociated  molecules  of  NH^OH  (resulting  from 
the  addition  of  NH^Cl),  the  equation  becomes :  — 

(NH,  4-  X  -y)  (OH  -  y)  _ 
NH4OH  +  y 

This  decreases  the  number  of  OH  ions  and  leaves  only  (OH  —  y) 
for  unit  volume.  It  is  owing  to  this  disappearance  of  OH  ions 
that  Mg(0H)2  is  not  precipitated,  —  too  little  molecular  Mg(0H)2 
being  formed  to  oversaturate  the  solution. 

(&)  Salts  of  Trivalent  Metals.  — When  NaOH,  NH^Cl,  and 
AlClg  are  brought  together  the  following  reactions  probably 
occur :  — 

AICI3  4-  3NaOH  =  Al(OH)3  -h  3NaCl ; 
NaOH-l-NH^Cl    =  NaCl  +  NII.OH. 


104  CHEMICAL  ANALYSIS 

Now  A1(0H)3,  unlike  Mg(0H)2  and  hydroxides  of  some  biva- 
lent metals,  is  a  very  weak  base  and,  consequently,  very  poorly 
dissociated.  Hence,"  when  an  excess  of  NH^Cl  is  added,  it  sup- 
presses only  the  OH  ions  of  NH^OH  —  not  those  of  A1(0H)3. 
Thus  NH4CI  not  only  does  not  affect  the  precipitated  hydroxide 
but  also  destroys  the  power  of  NH^OH  to  dissolve  it. 

Chromium  (salt  for  study,  chromium  sulphate,  Cr2(S04)3). 

1.  NH4OH Ogives  reactions  similar  to  1,  under  alumi- 
num. 

2.  NaOH  gives  reactions  similar  to  2,  under  aluminum. 

3.  BaCOg  gives  reactions  similar  to  3,  under  alumi- 
num, except  that  the  precipitation  requires  more  time 
for  its  completion. 

4.  (1014)28  gives  reactions  similar  to  4,  under  alumi- 
num. 

5.  Fusion  with  fusion  mixture  on  a  platinum  foil,  or 
with  sodium  dioxide,  Na202,  on  thick  silver  foil,  gives 
a  soluble  yellow  mass,  containing  sodium  chromate, 
NaaCrO^. 

6.  Na202  heated  with  a  solution  of  a  chromium  salt 
gives  yellow  Na2Cr04. 

7.  Borax  bead  with  both  oxidizing  and  reducing  flames 
gives  a  yellow-green  coloration  of  sodium  chromium 
metaborate,  NaQCr2(B02)i2- 

Reactions  5  and  6  illustrate  the  conversion  of  chromium  as  a 
base-producing  element  to  chromium  as  an  acid-producing  element. 

There  are  two  classes  of  chromium  compounds  derived  from 
the  two  oxides  CrgOg  and  CrOg. 

Chromic  oxide,  CrgOg,  is  basic  and  forms  salts  with  acids: 
CrgOg  -\-  6  HCl  =  2  CrClg  +  3  H2O.  By  oxidation  CrgOg  is  changed 
to  chromium  trioxide,  CrOg,  which  is  an  anhydride  and  forms 


METALS   OF  GROUP  III  105 

salts  with  bases:  CrOg  +  2NaOri  =  N'a2Cr04  +  HgO.  The  oxi- 
dation of  CrgOg  in  solution  may  be  accomplished  by  Na202  or 
by  hydrogen  dioxide,  HgOg.  CrOg  is  a  strong  oxidizing  agent 
and  is  easily  reduced  to  CrgOg  by  various  reagents,  namely, 
HgS,  SOg,  HCl,  and  many  organic  compounds.  If  HgS  is  passed 
through  an  acidified  (HCl)  solution  of  potassium  dichromate, 
KgCrgO^,  there  will  be  a  change  of  color  from  red  to  green : 
K2Cr207  +  3H2S  +  8  HCl  =  2CrCl3  +  2KC1  +  3S  +  7H2O. 

In  this  case  chromium  is  changed  from  the  acid  to  the  basic 
condition. 

CHEMISTRY   OF  BORAX   BEADS 

Borax  (sodium  tetraborate,  NagB^O^-  IOH2O),  like  NagCOg,  is 
both  an  inactive  flux,  as  in  normal  sodium  borate,  NagBOg,  and 
an  active  chemical  agent,  as  in  boric  acid,  HgBOg.  (See  Fusion, 
p.  43.)  When  borax  is  heated,  it  loses  its  water  of  crystallization 
and  fuses  to  a  clear  bead  on  the  platinum  wire.  If  a  metallic 
oxide  or  salt  is  fused  with  the  clear  bead,  a  double  borate  is 
formed,  which  is  often  colored.  The  reaction  can  be  easily 
understood  by  a  review  of  the  principal  hydroxyl  acids  of  boron. 
Water  unites  with  boric  oxide,  B2O3,  increasing  in  an  arithmet- 
ical progression  to  form  a  systematic  chain  of  polyboric  acids, 
of  which  the  following  are  important  in  this  connection :  — 

2  BgOg  +  HgO     =  H2B4O7,  dihydroxyl  tetraboric  acid ; 

2B20g  +  2H2O  =  4HBO2,  metaboric  acid; 

2  BgOg  +  3  HgO  =  HgB^Og,  hexahydroxyl  tetraboric  acid ; 

2  BgOg  +  4  H2O  =  2  H4B2O5,  diboric  acid ; 

2  BgOg  +  5  H2O  =  HjqB^Ou,  dekahydroxyl  tetraboric  acid ; 

2B20g  +  6H2O  =  4  HgBOg,  (normal)  orthoboric  acid. 

From  these  equations  it  can  be  seen  that  any  succeeding  acid 
in  the  list  can  be  formed  from  the  next  preceding  by  the  addition 
of  one  molecule  of  water ;  and  the  reverse  is  also  true,  that  any 
higher  acid  can  be  formed  from  the  next  lower. 

The  corresponding  salts  of  these  acids  are  formed  in  a 
similar  manner,  provided  the  factor  to  be  added  is  a  metallic 


106  CHEMICAL  ANALYSIS 

oxide  instead  of  water.     The  most  important  salts  are  given  for 
illustration :  — 

NagB^O,,  sodium  tetraborate  (borax) ; 

NagB^O^  +  NagO    =4]SraB02,         "        metaborate; 
NagB^O^  +  3  NagO  =  2  Na^B^O^,       "        diborate ; 
NagB^O^  +  5  NagO  =  4  NagBOg,        "        ortlioborate. 

This  further  explains  the  statement  above,  that  borax  is  both 
an  inactive  flux,  like  NagBOg,  and  an  active  chemical  agent,  like 

HgBOg. 

It  is  equivalent  to  four  molecules  of  boric  acid  with  one  mole- 
cule of  water  replaced  by  sodium  oxide,  and  the  other  five  mole- 
cules of  water  displaced  :  — 

4  HgBOg  +  NagO  =  Na^B^O^  +  6H2O. 

Thus  it  is  in  part  a  salt  and  in  part  an  anhydride. 
If  an  oxide  of  a  heavy  metal  is  substituted  for  NagO,  double 
borates  are  formed  :  — 

NagB^O^  +  CoO    =  N'a2Co(B02)4,  sodium  cobaltous  metaborate ; 
Na2B407  -I-  3  CoO  =  Na2Co3(B205)2,    "        cobaltous  diborate ; 
NagB^O^  +  5  CoO  =  Na2Co5(B03)4,      "        cobaltous  orthoborate. 

Or,  in  the  case  of  the  triad  element  chromium  :  — 

3Na2B407  +  CrgOg  =  NagCr2(B02)i2»  sodium  chromic  metaborate, 
etc. 

The  metaborate,  then,  is  the  first  product  formed  by  adding  a 
small  quantity  of  the  oxide  to  an  excess  of  borax,  while  the  ortho- 
borate is  the  last  formed  by  adding  a  larger  amount  of  the  oxide. 
What  the  actual  composition  of  a  given  bead,  made  without 
weighing  its  components,  may  be  can  only  be  determined  by 
a  quantitative  analysis.  Probably  every  bead  contains  more  or 
less  of  each  of  a  large  number  of  double  borates.  However,  as 
the  colors  can  be  seen  best  by  using  small  quantities  of  the 
oxides  with  a  large  excess  of  borax,  the  metaborates  predomi- 
nate, and  as  such  the  beads  are  usually  represented. 


METALS    OF   GROUP  III  107 

The  same  laws  which  apply  to  the  formation  of  polyborates 
apply  to  the  formation  of  polyphosphates  and  silicates.  Hence 
the  vast  number  of  natural  and  artificial  silicates  can  be  traced 
to  their  corresponding  acids. 

Ferrous  Iron,  Fe"  (salt  for  study,  ammonium  ferrous 
sulphate,  (NH4)2Fe"(S04)2). 

1.  NH4OH  and  NaOH  precipitate  ferrous  hydroxide, 
Fe (011)2,  which  oxidizes  quickly  to  brown  ferric 
hydroxide,  Fe(0H)3.  NH4CI  partly  prevents  the  pre- 
cipitation by  NH4OH,  and  partly  that  by  NaOH. 

2.  (NH4)2S  precipitates  black  ferrous  sulphide,  FeS. 

3.  K4Fe(CN)g  precipitates  white  potassium  ferrous 
ferrocyanide,  K2Fe"Fe(CN)g,  which  rapidly  oxidizes  to 
Prussian  blue. 

4.  K3Fe(CN)g  precipitates  TurnbuU's  blue,  ferrous 
ferricyanide,  Fe3''(Fe(CN)g)2. 

5.  Borax  bead  with  the  oxidizing  flame  gives  a  yellow 
coloration,  NaQFe2'"(B02)i2 ;  with  the  reducing  flame, 
a  green  coloration,  Na2Fe''(B02)4. 

Ferric  Iron,  Fe'"  (salt  for  study,  ferric  chloride,  FeClg). 

1.  NH4OH  and  NaOH  precipitate  brown  ferric  hydrox- 
ide, Fe(0H)3,  insoluble  in  excess  of  reagents.  NH^Cl 
does  not  prevent  the  precipitation  either  by  NH^OH  or 
by  NaOH. 

2.  BaCOg  precipitates  a  brown  basic  salt,  Fe20(C03)2. 

3.  H2S  reduces  ferric  to  ferrous  salts  and  precipi- 
tates free  sulphur:  — 

2  FeCl3  +  H2S  =  2FeCl2  +  2HC1  +  S. 


108  CHEMICAL   ANALYSIS 

The  precipitate  of  sulphur  formed  iu  this  reaction  is  sometimes 
mistaken  for  the  sulphides  of  certain  members  of  Group  II. 
This  confusion  will  not  occur  if  it  is  recalled  that  the  members  of 
Group  II  are  all  colored,  whereas  a  precipitate  of  finely  divided 
sulphur  is  white. 

^     4.   (NH4)2S  reduces  ferric  to  ferrous  salts  and  precipi- 
tates ferrous  sulphide :  — 

2FeCl3  +  3  (NH4)2S  =  2FeS  +  GNH^Cl  +  S. 

5.  K4Fe(CN)g  precipitates  Prussian  blue,  ferric  ferro- 
cyanide,  Fe/"(Fe(CN)e)3. 

6.  K3Fe(CN)g  gives  a  brown  coloration. 

7.  KCNS  gives  a  deep-red  coloration,  ferric  sulpho- 
cyanate,  Fe(CNS)3.i 

8.  Borax  bead  gives  the  same  results  as  with  ferrous 
salts. 

PKOCESS   OF   SEPARATION 

The  separation  of  the  members  of  this  group  is  based 
upon  the  facts  that  Cr(0H)3  is  oxidized  to  soluble 
NagCrO^  by  means  of  Na202  or  by  fusion  with  the 
mixture  of  Na2C03  and  KNO3,  and  that  A1(0H)3  is 
soluble  in  an  excess  of  NaOH.  Two  methods  of  mak- 
ing the  separation  are  given,  of  which  the  first  and 
simpler  is  to  be  employed  in  the  absence  of  phosphoric, 
boric,  silicic,  and  hydrofluoric  acids;  whereas  the  sec- 
ond and  more  complicated  method  is  to  be  followed  in 
the  presence  of  these  bodies.^ 

Boil  off  all  traces  of  H2S,  testing  for  its  removal  by 
holding  above  the  liquid  a  strip  of  paper  moistened 
with  AgN03   or   Pb(N03)2.     It   must  be   di'iven  off 


METALS    OF  GROUP  III  l09 

completely;  since,  if  it  were  allowed  to  remain,  the 
members  of  Group  IV  would  be  precipitated  out  of  due 
course  upon  the  addition  of  NH4OH,  the  precipitant 
for  the  members  of  Group  III.  Next  test  for  iron  by 
adding  K3Fe(CN)g  to  a  small  portion  of  the  solution; 
and  in  case  it  is  present  add  a  few  drops  of  IINO3  and 
boil  until  the  reaction  for  ferrous  compounds  disappears. 
The  oxidation  of  iron  to  the  ferric  state  at  this  point 
is  necessary,  since,  if  left  in  the  ferrous  state,  it 
would  not  be  precipitated  by  NH^OH  in  presence 
of  NH4CI. 

The  solution  should  now  be  tested  for  oxalic  acid 
or  other  organic  matter  by  evaporating  a  small  portion 
of  the  solution  to  dryness,  and  heating  the  residue  in 
a  closed  tube  connected  with  a  small  rubber  delivery 
tube,  through  which  any  gas  that  may  be  evolved  can 
be  conducted  into  lime  water.  A  charred  residue  in 
the  closed  tube  and  a  white  precipitate  in  the  lime 
water  indicate  the  presence  of  organic  compounds. ^ 
If  such  are  found,  evaporate  the  whole  solution  to 
dryness  and  heat  the  residue  with  the  addition  of  a 
few  drops  of  sulphuric  acid,  until  the  organic  matter 
is   thoroughly  decomposed.     Cover   the    residue  with 

1  In  testing  for  organic  matter  blackening  is  not  conclusive.  Many  inor- 
ganic salts,  among  them  those  of  iron,  nickel,  cobalt,  and  manganese, 
blacken  when  heated.  Furthermore,  a  failure  to  blacken  is  not  an  evidence 
of  the  absence  of  organic  compounds,  since  some  of  them  which  contain  a 
large  per  cent  of  oxygen  —  oxalates  in  particular  —  do  not  char,  but  give 
off  all  their  carbon  as  oxides.  The  lime-water  test,  too,  is  not  absolute, 
though  more  reliable  than  that  by  charring.  Compounds  evolving  oxides 
of  sulphur  also  whiten  lime  water.  But  both  tests  are  good  signs ;  and  as 
gentle  ignition  is  also  the  means  of  eliminating  silicic  acid,  it  is  best  to  evap- 
orate and  ignite  the  solution  even  when  the  presence  of  organic  acids  is 
doubtful.     V  "  "' 


110  CHEMICAL   ANALYSIS 

concentrated  HCl,  evaporated  almost  to  dryness,  add 
water  and  a  few  drops  of  concentrated  HNO3,  and  boil. 
Filter  the  solution  from  the  separated  carbon  and  silica. 
Phosphoric  acid^  must  next  be  tested  for  by  warm- 
ing a  small  portion  of  the  filtrate  with  an  excess  of 
HNO3  solution  of  (NH4)2Mo04.  Should  it  be  present, 
barium  must  also  be  tested  for  at  this  stage  by  making 
a  small  portion  alkaline  with  NH^OH,  acidifying  with 
HC2H3O2,  and  testing  for  barium  with  K2Cr04.  Next 
add  NH^Cl,^  boil,  and  add  NH^OH  till  its  odor  persists. 
Filter  quickly  while  hot.  Reserve  filtrate  (a)  for  subse- 
quent groups.  Redissolve  residue  (a)  in  least  quantity 
of  HCl,  nearly  neutralize  with  Na2C03,  transfer  to  a 
stoppered  flask,  and  add  when  cold  a  large  excess  of 
suspended  BaCOg.^  Shake  from  time  to  time,  and  filter 
after  15  minutes.  If  phosphoric  acid  is  present,  combine 
filtrates  (a)  and  (h)^ ;  if  absent,  test  filtrate  {b)  for  man- 
ganese by  evaporating  to  dryness  and  fusing  with  NagCOg 
and  KNOg.  Residue  (h)  may  consist  of  basic  salts  of 
Group  III,  if  phosphoric  acid  is  absent ;  or,  if  present,  it 
may  also  contain  phosphates  of  metals  of  Groups  III, 
IV,  V,  and  of  magnesium. 

In  Absence  of  Phosphates 
Method  1.  Thoroughly  wash  the  residue  {b)  and  trans- 
fer to  a  test-tube.  Add  a  small  quantity  of  water  and 
some  bits  of  Na202,  and  boil  till  effervescence  ceases. 
Filter  and  wash.  The  residue  {c)  may  be  brown  Fe(OH)g, 
whose  identity  can  be  confirmed  by  dissolving  in  dilute 
HCl  and  testing  with  K4Fe(CN)6.  The  filtrate  {e)  may 
contain  yellow  NagCrO^  and  NagAlOg. 


METALS    OF   GROUP   III  111 

The  following  equations  explain  the  formation  of  the 
soluble  salts :  — 

3  NaaOa  +  2  Cr(OH)3  =  2  Na2Cr04+  2  H2O  +  2  NaOH  ; 
6  NaOH  +  2  A1(0H)3=  2  Na3A103+  6  II2O. 

Divide  the  filtrate  (c)  into  two  parts.  Acidify  the 
one  with  HCgHgOg  and  test  for  chromium  by  adding 
Pb(C2H302)2.  Acidify  the  other  part  with  dilute  HCl, 
and  while  boiling  test  for  aluminum  by  adding  an 
excess  of  NH^OH,  —  or  to  the  second  part  of  the 
filtrate  add  some  NH^Cl  and  boil.  After  cooling,  a 
gelatinous  precipitate  confirms  presence  of  aluminum. 

Method  2.  Dry  residue  (h)  and  fuse  in  a  platinum 
crucible  or  foil  with  an  excess  of  fusion  mixture.  The 
cooled  mass  is  triturated  in  a  mortar,  —  preferably  a 
glass  one,  —  is  digested  with  water  for  fifteen  minutes, 
and  then  is  filtered.  The  residue  {c)  is  tested  for  iron, 
and  the  filtrate  {c)  for  chromium  and  aluminum,  as  in 
Method  1.  In  case  of  doubt  test  the  solution  for 
aluminum  with  the  spectroscope. 

In  Presence  of  Phosphates 

Dissolve  residue  {h)  in  a  little  HCl,  nearly  neutralize 
by  cautious  addition  of  Na2C03,  add  NaC2H302  and 
HC2H3O2,  and  boil  and  filter.  The  filtrate  (c)  may  con- 
tain phosphates  of  the  metals  of  Groups  IV  and  V, 
and  of  magnesium.  The  residue  {c)  may  consist  of  the 
phosphates  of  aluminum,  chromium,  and  iron.  To  the 
filtrate  (c)  add  dilute  FeClg,  drop  by  drop,  until  a 
red  coloration  appears.     At  this  point,  precipitation  of 


112  CHEMICAL  ANALYSIS 

FePO^  is  completed.  The  red  color  indicates  the 
formation  of  Fe(C2Hg02)3,  an  excess  of  which  would 
redissolve  the  FePO^.  Hence  the  solution  should  now 
be  boiled  in  order  to  change  any  excess  of  Fe(C2H302)3 
to  an  insoluble  basic  acetate,  FeO(C2H302).  It  is  neces- 
sary to  filter  the  mixture  with  the  pump  while  hot,  as 
FeO(C2H302)  redissolves  to  Fe(C2H302)3  on  cooling. 
The  filtrate  (d)  may  contain  the  chlorides  of  Groups  IV 
and  V,  and  of  magnesium.  This  filtrate  should  be 
combined  with  filtrate  (a)  and  afterwards  tested  for  sub- 
sequent groups.  The  residue  (<?),  possibly  consisting  of 
FeP04  and  some  FeO(C2H302),  should  be  rejected. 

The  separation  of  the  members  of  Group  III,  which 
may  be  present  as  phosphates  in  residue  (J),  can  now 
be  made  without  difficulty  by  either  of  the  procedures, 
Methods  1  and  2. 

The  following  further  explanations  may  be  given  concerning 
the  method  of  analysis  in  presence  of  phosphates :  — 

It  will  be  seen,  on  reference  to  the  Table  of  Solubilities  given 
on  p.  34,  that  the  phosphates  of  Mg,  Ba,  Sr,  Ca,  Co,  Ni,  Mn,  and 
Zn  are  all  more  or  less  insoluble  in  water  or  alkaline  solutions. 
If  present  in  an  acid  solution,  they  therefore  will  be  thrown  out 
upon  neutralization  ;  and  if  this  solution  be  under  examination 
for  the  members  of  Group  III,  they  will  separate  simultaneously 
with  the  hydroxides  of  this  group.  Hence  the  necessity  of  follow- 
ing the  modified  method  of  analysis  in  their  presence.  They  play 
so  important  a  part  in  the  separation  of  the  members  of  Group  III 
that  the  student  is  advised  to  perform  in  advance  the  exercises 
with  phosphoric  acid,  given  on  p.  135,  in  order  that  he  may 
have  some  practical  knowledge  of  their  reactions. 

In  removing  phosphoric  acid  after  almost  neutralizing  the  HCl 
solution  with  NagCOg,  NaCgHaOg  is  added  to  destroy  the  solvent 
effect  of  HCl,  which  dissolves  all  of  the  phosphates.     HCgHgOg, 


METALS   OF  GROUP  III  113 

which  is  set  free  from  NaCgHgOg  by  metathesis  with  HCl,  dissolves 
only  the  phosphates  of  Groups  IV  and  Y,  and  magnesium,  —  not 
those  of  Group  III.  But,  though  HCgHgOg  is  set  free,  its  acid 
effects  are  greatly  weakened  by  the  presence  of  an  excess  of  a  salt 
having  an  ion  in  common  with  it ;  and  for  this  reason  some  free 
HC2H3O2  must  be  added  with  N'aCgHgOg,  though  a  large  excess 
of  the  acid  should  be  avoided,  as  CrPO^  is  sparingly  soluble  in 
it.     (See  Question  5,  Theory  of  Solutions,  p.  22.) 


CHAPTER    X 

METALS   OF   GROUP   IV:    ZINC,   MANGANESE,   COBALT, 
AND   NICKEL 

CiiAKACTEUiSTic  :  Insolubility  of  the  sulphides  in  alkaline  solu- 
tion. 

Guoup  IvKAdENT  :  Ammonium  sulphide  in  presence  of  ammo- 
nium hydroxide  and  ammonium  chloride. 

REACTIONS 

Zinc  (salt  for  study,  zinc  sulphate,  ZnS04). 

1.  (^3:4)28  precipitates  white  zinc  sulphide,  ZnS, 
soluble  in  strong  acids ;  insoluble  in  alkalies  and 
HC2H3O2.  Tlie  precipitation  is  hastened  in  dilute 
solutions  by  presence  of  NH^Cl. 

2.  NH4OH  and  NaOH  precipitate  white  zinc  hydrox- 
ide, Zn(0H)2,  soluble  in  excess  of  reagent,  forming 
ammonium  zincate,  (NH4)2Zn02,^  or  sodium  zincate, 
Na2Zn02.     NH^Cl  solutions  dissolve  Zn(0H)2. 

3.  Reducing  flame  on  charcoal  with  NagCOg  gives  a 
yellow  incrustation  of  zinc  oxide,  ZnO,  turning  white  on 
cooling.  If  this  coating  is  moistened  with  cobalt  nitrate 
and  again  heated  with  the  blowpipe,  a  green  coloration 
will  appear,  due  to  a  double  oxide  of  zinc  and  cobalt. 

Manganese  (salt  for  study,  manganese  sulphate,  MnS04). 

1.  (NH4)2S  precipitates  pink^  manganous  sulphide, 
MnS,  soluble  in  acids,  including  HC2H3O2 ;  insoluble 
in  alkalies.     NH^Cl  assists  precipitation. 

114 


METALS   OF  GROUP  IV  115 

2.  NH4OH  and  NaOH  precipitate  manganous  hydrox- 
ide, Mn{0H)2,  which  oxidizes  to  brown  hydrated  man- 
ganic oxide,  Mn202(OH)2.  NH^Cl  partly  redissolves 
Mn(0H)2  if  precipitated  by  NH^OH,  but  only  partially 
if  precipitated  by  NaOH. 

3.  Fusion  with  Na2C03  and  KNO3  on  a  platinum  foil 
gives  a  soluble  green  mass  of  sodium  manganate, 
NagMnO^,  changing  to  red  sodium  permanganate, 
Na2Mn20g,  on  heating  or  acidifying.  This  reaction  is 
also  produced  by  adding  Na202  to  the  solution  and 
warming  till  effervescence  ceases. 

4.  Boiling  with  Pb02  and  H2SO4  gives  a  deep-red 
coloration  of  permanganic  acid,  HgMugOg. 

5.  Borax  bead  in  the  oxidizing  flame  gives  an  amethyst 
coloration,  sodium  manganic  metaborate,  NagMn2(B02)i2- 
This  color  is  destroyed  })y  the  reducing  flame,  due  to  the 
formation  of  manganous  metaborate. 

6.  Spectrum  (see  Special  Method  for  Manganese,  p.  63). 

CHEMISTRY   OF   MANGANESE    COMPOUNDS 

There  are  six  important  oxides  of  manganese  having  the  follow- 
ing formulae  :  MnO,  Mn304,  MngOs,  MnOa,  MnOg,  MngO^.  Of 
these,  the  first,  second,  and  third  are  basic,  the  first  forming  man- 
ganous salts  with  acids,  and  the  third,  manganic  salts.  The  fourth 
oxide  is  both  basic  and  acid,  forming  manganous  salts  with  loss  of 
oxygen  and  also  forming  unstable  manganites  with  more  basic 
oxides.  The  fifth  and  sixth  are  acid  oxides  and  produce,  respec- 
tively, manganates  and  permanganates.  The  ready  conversion  of 
the  lower  basic  oxides  to  the  higher  acid  oxides  by  oxidation  is  of 
importance  in  qualitative  analysis,  on  account  of  the  solubility  and 
characteristic  colors  of  the  salts  of  the  higher  oxides.  Reactions 
3,  4,  and  5  illustrate  these  effects. 


116  CHEMICAL  ANALYSIS 

The  higher  oxides  are  also  easily  reduced  to  the  lower,  and 
for  this  reason  are  extensively  used  in  quantitative  analysis. 

The  following  equations  illustrate  a  few  of  the  many  cases  of 
oxidation  depending  on  the  reduction  of  the  higher  oxides  of 
manganese  :  — 

(a)  Oxidation  of  oxalic  acid :  — 

5H2C2O4  +  K2Mn208  +  3H2SO4  =  SHgO  +  IOCO2  +  2MnS04 
+  K2SO4; 

(b)  Oxidation  of  hydrochloric  acid :  — 

KgMngOg  +  16  HCl  =  2  KCl  +  2  MnCla  +  8  HgO  +  10  CI ; 

(c)  Oxidation  of  ferrous  to  ferric  salts :  — 

lOFeSO^  +  K2Mn208  +  8  H2SO4  =  5  Fe^(SO^)^  +  K2SO4 
+  2MnS04  +  8H20. 

In  all  these  cases  the  reduction  of  KaMugOg  is  accompanied  by 
the  destruction  of  the  brilliant  red  color  of  the  salt. 


Cobalt  (salt  for  study,  cobalt  nitrate,  Co(N03)2). 

1.  (NH4)2S  precipitates  black  cobaltous  sulphide, 
CoS,  insoluble  in  cold  dilute  HCl ;  soluble  in  HNO3 
and  in  aqua  regia.  The  presence  of  NH^Cl  aids  the 
precipitation  of  CoS. 

2.  NH4OH  precipitates  blue  cobaltous  hydroxide, 
Co(OH)2,  soluble  to  a  brown  fluid  in  excess  of  re- 
agent.    NH4CI  hinders  the  precipitation. 

3.  NaOH  precipitates  blue  cobaltous  hydroxide, 
Co(OH)2,  not  soluble  in  excess  of  reagent.  NH^Cl 
prevents  precipitation. 

4.  KCN  precipitates  light-brown  cobaltous  cyanide, 
Co(CN)2,  soluble  in  excess  of  reagent  with  formation 
of  potassium  cobaltous  cyanide,  K4Co(CN)g.  This 
compound  is  decomposed  by  acids,  in  the  absence  of 


METALS  OF  GROUP  IV  117 

an  excess  of  KCN,  with  reprecipitation  of  Co(CN)2. 
If  K4Co(CN)g  is  oxidized  with  a  mixture  of  NaOH  and 
bromine  water,  containing  NaBrO,  it  changes  to  the 
stable  -10  salt,  K3Co(CN)g,  which  will  yield  no  precipi- 
tate with  acids. 

5.  KNO2  added  to  the  solution  strongly  acidified  with 
HC2H3O2  gives  a  yellow  precipitate,  potassium  cobaltic 
nitrite,  K3Co(N02)6.  The  precipitate  will  separate 
after  some  hours  in  a  warm  place. 

K3Co(N02)6  is  readily  broken  down  by  strong  acids  and  alkalies, 
but  .is  insoluble  in  HCgHgOg  and  in  a  solution  of  KNOg.  Hence, 
if  strong  acid  is  present,  it  is  necessary  to  neutralize  it  with 
NagCOg  and  then  to  add  an  excess  of  HCgHgOg  and  KNOg. 

6.  Borax  bead  with  both  the  oxidizing  and  reducing 
flames  gives  a  blue  coloration  of  sodium  cobaltous 
metaborate,  Na2Co(B02)4. 

Nickel  (salt  for  study,  nickel  chloride,  NiCl2). 

1.  (1014)28  precipitates  black  nickel  sulphide,  NiS, 
sparingly  soluble  in  excess  of  reagent ;  soluble  in  hot 
HNO3  or  aqua  regia  ;  almost  insoluble  in  dilute  HCl. 
The  presence  of  NH^Cl  aids  the  precipitation  of  NiS. 

2.  NH4OH  precipitates  greenish  nickelous  hydroxide, 
Ni(0H)2,  soluble  in  excess  of  reagent,  forming  a  blue 
fluid.     NH4CI  prevents  the  precipitation. 

3.  NaOH  precipitates  Ni(0H)2,  insoluble  in  excess  of 
reagent,  but  soluble  in  NH^Cl. 

4.  KCN  precipitates  greenish  nickelous  cyanide, 
Ni(CN)2,  soluble  in  excess  of  reagent,  forming  potas- 
sium nickelous  cyanide,  K2Ni(CN)4.     NaBrO  does  not 


118  CHEMICAL  ANALYSIS 

oxidize  K2^i(^^)4  ^^  ^^  corresponding  -ic  salt,  but  to 
black  nickelic  hydroxide,  Ni(0H)3. 

5.  KNO2  produces  no  precipitate  with  nickel  salts. 

6.  Borax  bead  in  the  oxidizing  flame  gives  yellow 
NagNi  (602)4;  ^^  ^^  reducing  flame  it  gives  gray 
metallic  nickel. 

Why  concentrated  IICl  does  not  dissolve  CoS  and  NiS  has  not 
been  satisfactorily  explained.  It  would  seem  reasonable  that  as 
II2S  does  not  precipitate  the  sulphides  of  cobalt  and  nickel  from 
a  solution  of  their  chlorides,  their  sulphides  ghould  be  soluble  in 
HCl.  It  has  been  surmised  that  immediately  after  precipitation 
the  sulphides  undergo  polymerization^  i.e.,  a  locking  together  of 
several  of  their  molecules,  to  form  very  insoluble  compounds, 
(CoS)x  and  (NiS)^. 

PROCESS   OF  SEPARATION 

The  separation  of  the  members  of  this  group  is  based 
upon  the  facts  that  Zn(0H)2  is  soluble  in  excess  of 
NaOH;  that  MnS  is  soluble  in  HC2H3O2 ;  and  that 
the  borax  bead,  KNO2,  and  KCN  with  NaBrO  give  dis- 
tinctive reactions  with  cobalt  and  nickel  salts.  The 
process  of  separation  is   as  follows :  — 

Boil  oif  excess  of  NH^OH,  add  NH^Cl,  and  then 
(NH4)2S  in  moderate  excess.  Filter  and  wash^  thor- 
oughly, rejecting  the  washings.  The  residue  (a)  may 
consist  of  the  sulphides  of  the  group.  The  filtrate  {a) 
may  contain  members  of  Groups  V  and  VI,  and  pos- 
sibly Mn,  which  separates  slowly  from  dilute  solutions. 
MnS  oxidizes  readily  to  brown  Mn202(OH)2,  —  espe- 
cially if  in  dilute  solutions,  or  if  the  residue  is  exposed 
to  the  air  for  a  short  while,  —  and  this  may  appear  as 


METALS   OF  GROUP  IV  119 

a  brown  precipitate  ^  in  the  filtrate  from  the  sulphides 
of  Group  IV.  Therefore  set  this  filtrate  aside,  and 
if  such  a  precipitate  appears,  filter  and  preserve  the 
filtrate  to  be  tested  for  Groups  V  and  VI. 

Dissolve  the  residue  (a)  in  boiling  dilute  HCl,  to 
which  a  small  crystal  of  KCIO3  is  added.  Continue 
the  boiling  till  all  free  chlorine  is  expelled ;  then  add 
while  stirring  an  excess  of  NaOH.  After  cooling,  filter. 
The  filtrate  (b)  may  contain  NagZnOg,  and  the  residue  (b) 
may  consist  of  Mn(0H)2,  Co(OH)2,  and  Ni(0H)2.  Pass 
H2S  through  the  filtrate  (b).  A  white  precipitate  con- 
firms the  presence  of  zinc. 

The  washed  residue  (b)  is  dissolved  in  a  small 
quantity  of  hot  HCl.  After  nearly  neutralizing  with 
NH4OH,  some  NH^CgHgOg  is  added  and  IT2S  is  passed 
till  the  precipitation  is  completed.  Filter.  The  resi- 
due (c)  may  consist  of  CoS  and  NiS.  The  filtrate  (c) 
may  contain  Mn(C2H302)2,  and  is  to  be  concentrated 
to  a  small  bulk  and  tested  in  either  of  the  following 
ways :  — 

First:  Add  NH^OH,  NH^Cl,  and  (NH4)2S.  A  yel- 
low or  green  precipitate,  appearing  after  some  time, 
confirms  MnS. 

Second:  Add  a  solution  of  Na2C03,  dissolve  the 
white  precipitate  in  HCl,  and  add  NH^OH,  NH^Cl, 
and  (NH4)2S.  In  case  of  doubt,  test  the  solution  for 
manganese  with  the  spectroscope. 

Test  the  residue  (c)  with  the  borax  bead  for  cobalt.  - 

Dissolve  the  residue  (c)  in  a  very  little  aqua  regin, 
evaporate  off  chlorine  and  excess  of  acids,  and  divide 
into  two  parts. 


120  CHEMICAL  ANALYSIS 

To  one  portion  add  Na2C03  till  alkaline,  then 
HC2H3O2  in  excess,  and  finally  add  some  solid  KNOg. 
After  standing  twenty-four  hours  in .  a  warm  place,  a 
yellow  precipitate  confirms  K3Co(N02)6-  Filter  and 
add  excess  of  NaOH,  which  will  precipitate  Ni(0H)2 
if  nickel  is  present. 

Make  the  second  part  neutral  by  adding  NaOH,  test- 
ing with  litmus  paper  ;  then  add  KCN  until  the  yellow 
precipitate  redissolves.  Digest  with  constant  stirring 
for  10  minutes,  or  till  the  dark  color  disappears.  Filter 
into  a  large  test-tube,  and  add  to  the  filtrate  an  equal 
bulk  of  NaOH  and  sufficient  bromine  water  to  produce 
a  permanent  red  color.  A  black  precipitate  on  gently 
warming  confirms  nickel. 


CHAPTER   XI 

METALS  OF  GROUP  V  :  BARIUM,  STRONTIUM,  AND  CALCIUM 

Characteristic  :  Insolubility  of  the  carbonates  in  alkaline 
solution. 

Group  Reagent  :  Ammonium  carbonate  in  presence  of  ammo- 
nium hydroxide  and  ammonium  chloride. 

REACTIONS 
Barium  (salt  for  study,  barium  nitrate,  Ba(N03)2). 

1.  (NH4)2C03  precipitates  white  barium  carbonate, 
BaCOg,  soluble  in  acids  (except  H2SO4)  and  in  acid 
ammonium  carbonate,  H(NH4)C03.  As  (NH4)2C03 
easily  dissociates  into  H(NH4)C03  and  NH3,  it  is 
necessary  to  add  NH^OH  before  (NH4)2C03.  NH^Cl 
solutions  dissolve  BaC03  slightly,  —  especially,  while 
boiling,  —  giving  off  NH3  and  CO2. 

2.  H2SO4  (also  CaSO^  and  SrSO^)  precipitates  white 
barium  sulphate,  BaSO^,  insoluble  in  dilute  acids  and 
alkalies;  somewhat  soluble  in  hot  concentrated  acids. 

3.  (NH4)2C204  precipitates  white  barium  oxalate, 
BaCgO^,  soluble  in  acids,  including  HC2H3O2. 

4.  K2Cr04  precipitates  yellow  barium  chromate, 
BaCrO^,  soluble  in  HCl  and  HNO3 ;  somewhat  solu- 
ble in  HC2H3O2  and  in  NH^Cl;  insoluble  in  potas- 
sium dichromate,  K2Cr20y,  which  latter  salt  is  formed 
by  the  action  of  acids  on  K2Cr04  :  — 

2K2Cr04  +  2  HCl  =  K2Cr207  +  2  KCl  +  H2O. 
121 


122  CHEMICAL  ANALYSIS 

Hence,  in  acid  solutions  of  barium  salts,  KgCrO^ 
should  be  added  in  excess  to  neutralize  the  acids. 

5.  Ether-alcohol  does  not  dissolve  barium  nitrate, 
Ba(N03)2.  As  Ba(N03)2  is  soluble  in  water,  it  is 
necessary  to  conduct  the  experiment  with  a  perfectly 
dry  salt. 

6.  Heated  on  a  platinum  wire  in  the  flame  of  the 
Bunsen  lamp,  barium  salts  give  a  yellow-green  color, 
which,  seen  through  a  cobalt  glass,  appears  blue-green. 

7.  Spectrum  (see  Table  VII,  p.  58). 

Strontium  (salt  for  study,  strontium  nitrate,  Sr(N03)2). 

1.  (NH4)2C03  precipitates  white  strontium  carbonate, 
SrCOg,  which  for  the  most  part  behaves  towards  rea- 
gents like  BaCOg.  It  is  less  soluble  in  NH^Cl  than 
BaCOg. 

2.  H2SO4  (also  CaSO^)  precipitates  white  strontium 
sulphate,  SrSO^,  sparingly  soluble  in  water,  but  more 
soluble  in  HCl  and  HNOg ;  insoluble  in  (NH4)2S04. 

3.  (NH4)2C204  gives  a  reaction  similar  to  3,  under 
barium,  except  that  SrC204  is  almost  insoluble  in 
HC2H3O2. 

4.  KjCrO^  does  not  precipitate  yellow  strontium  chro- 
mate,  SrCrO^,  except  on  long-standing  and  in  concen- 
trated neutral  solutions.  SrCr04  is  quite  soluble  in 
HC2Hg02. 

5.  Ether-alcohol  does  not  dissolve  strontium  nitrate, 
Sr(N03)2,  but  as  it  is  very  soluble  in  water,  it  is  essen- 
tial to  conduct  the  experiment  with  a  freshly  heated 
anhydrous  salt. 


A 


METALS   OF  GROUP    V  123 

6.  Heated  in  a  non-luminous  flame  on  a  platinum 
wire,  strontium  salts  give  a  deep  red  color,  which,  seen 
through  a  blue  glass,  appears  blue. 

7.  Spectrum  (see  Table  VII,  p.  58). 

Calcium  (salt  for  study,  calcium  nitrate,  Ca(N03)2). 

1.  (NH4)2C03  gives  a  reaction  similar  to  1,  under 
barium,  except  that  CaCOg  is  more  soluble  than  BaCOg 
in  NH4CI  and  in  H2SO4. 

2.  H2SO4  precipitates  white  calcium  sulphate,  CaSO^, 
sparingly  soluble  in  water  and  acids;  insoluble  in 
alcohol.  In  this  latter  experiment  it  is  necessary 
that  the  calcium  salt  be  in  a  concentrated  solution. 

3.  (NB.^)2C^0^  precipitates  white  calcium  oxalate,  sol- 
uble in  strong  acid;  insoluble  in  HC2H3O2. 

4.  K2Cr04  gives  no  precipitate  if  the  solution  is  acidi- 
fied with  HC2Hg02. 

5.  Ether-alcohol  dissolves  calcium  nitrate,  Ca(N0g)2. 

6.  Heated  in  a  non-luminous  flame  on  a  platinum  wire, 
calcium  salts  give  a  pale  red  color,  which  appears  green- 
gray  if  seen  through  a  blue  glass. 

7.  Spectrum  (see  Table  VII,  p.  58). 

PROCESS   OF   SEPARATION 
The  separation  of  the  members  of  this  group  may  be 
made  by  three  methods,  which  are  based  on  the  follow- 
ing phenomena :  — 

Method  1.     (a)  Insolubility  of  Ba(N03)2  and  Sr(N03)2 
in  ether-alcohol; 
(b)  Insolubility  of  BaCrO^  in  KgCrgO^  solu- 
tion. 


124  CHEMICAL  ANALYSIS 

Method  2.     (a)  Insolubility  of  BaCrO^  in  KgCigO^  solu- 
tion ; 
(b)  Insolubility  of  SrSO^  in  (NH4)2S04  solu- 
tion. 

Method  3.     Differences  between  the  spectra  of  barium, 
strontium,  and  calcium  salts. 

The  solution  is  first  made  alkaline  with  NH^OH, 
NH4CI  is  added  in  moderate  quantity,  and  finally 
(NH4)2C03  solution  is  added  till  the  precipitation  is 
completed.  The  mixture  should  be  warmed  (not  boiled), 
filtered,  and  washed  with  ammoniated  water.  The 
filtrate  {a)  may  contain  members  of  Group  VI.  The 
residue  {a)  may  consist  of  BaCOg,  SrCOg,  and  CaCOg. 

Method  1.  —  Dissolve  the  residue  (a)  in  an  evaporat- 
ing dish  with  a  minimum  of  dilute  HNO3,  and  evapo- 
rate carefully  to  dryness.  Remove  to  an  iron  plate  (a 
small  clean  sand  bath  may  be  used)  and  heat  to  dull 
redness,  till  all  traces  of  moisture  are  expelled,  tested 
by  holding  a  cold  dry  watch-glass  or  beaker  over  the 
dish.  After  cooling,  quickly  triturate  the  mass  in  a 
dry  mortar  with  about  10  c.c.  ether-alcohol.  Transfer 
to  a  small  dry  flask  and  shake  the  corked  flask  at  inter- 
vals. After  about  an  hour  filter  through  a  dry  paper 
(filtrate  (h))  and  wash  the  residue  (h)  with  a  little  ether- 
alcohol  till  the  drippings  show  no  cloudiness  with  dilute 
H2SO4.  Dissolve  the  residue  (5),  possibly  consisting 
of  Ba(N0g)2  and  Sr(N03)2,  in  warm  water  acidified 
with  a  few  drops  of  HC2H3O2,  and  while  boiling  add 
K2Cr04  till  the  solution  ceases  to  smell  of  HC2H3O2. 
A  yellow  precipitate  confirms  BaCrO^.  Filter,  and  to 
the  filtrate  {c)  add  NH4OH  and  (NH4)2C03.  A  white 
precipitate   confirms    SrCOg.     To   the  filtrate   {b)  add 


METALS   OF  GROUP   V  125 

some    dilute    H2SO4.       A    white    precipitate   confirms 

CaSO^. 

In  the  separation  of  the  nitrates  of  the  group  with  ether-alco- 
hol, not  only  the  external  moisture  but  also  the  water  of  crystal- 
lization of  the  nitrates  must  be  expelled.  The  temperature  may 
reach  180°  and  not  injure  the  salts,  but  above  that  degree  the 
nitrates  are  dissociated  into  the  insoluble  oxides. 

Method  2. — Dissolve  the  residue  {a)  in  a  small  amount 
of  dilute  HC2H3O2  and  add  Ys.^yO^  till  the  solution 
ceases  to  smell  of  HC2H3O2.  A  yellow  precipitate, 
residue  (J),  confirms  BaCrO^.  Filter.  Add  to  the 
filtrate  (b)  a  concentrated  solution  of  (NH4)2S04,  and 
boil.  Filter.  The  residue  {c)  indicates  SrS04,  which  can 
be  confirmed  with  the  flame.  Add  to  the  filtrate  {c)  some 
i^^^^^^^.     A  white  precipitate  confirms  CaCgO^. 

Method  3.  —  Dissolve  the  residue  (a)  in  the  least  pos- 
sible amount  of  HCl  and  evaporate  to  a  thin  paste. 
Test  with  the  spectroscope  for  barium,  strontium,  and 
calcium. 

Magnesium,  in  many  respects,  behaves  like  the  three  other 
alkali-earth  metals  —  barium,  strontium,  and  calcium.  In  one 
respect,  however,  magnesium  differs  very  materially  from  the 
other  kindred  metals.  Its  salts  readily  dissolve  in  ammonium 
salts,  but  only  to  a  very  Umited  degree  do  ammonium  salts  affect 
the  solubility  of  the  salts  of  barium,  strontium,  and  calcium. 
Hence  NH^Cl  must  be  added  with  (^'114)2003  to  prevent  the 
precipitation  of  MgCOg,  but  an  excess  of  the  reagent  must  be 
avoided  lest  it  also  dissolve  the  other  carbonates.  Heating  is 
usually  conducive  to  precipitation  ;  but  in  this  case  great  heat 
dissociates  (NH4)2C03  into  NH3  and  H(NH4)2C03,  which  latter 
acid  salt,  being  poorly  ionized,  does  not  precipitate  the  carbon- 
ates of  the  group.  Hence  the  solution  must  be  warmed,  but  not 
boiled. 


CHAPTER   XII 

METALS  OF  GROUP  VI:    MAGNESIUM,   AMMONIUM, 
POTASSIUM,  AND   SODIUM 

No  group  characteristic.    No  g^oup  reagent. 

REACTIONS 
Magnesium  (salt  for  study,  magnesium  chloride,  MgCl2). 

1.  HNa2P04  precipitates  white  crystalline  acid  magne- 
sium phosphate,  HMgPO^,  very  soluble  in  acids ;  some- 
what soluble  in  water.  If  NH^Cl  and  NH^OH  are 
added  before  HNagPO^,  the  more  insoluble  (NH4)MgP04 
will  be  precipitated.  It  is  soluble  in  acids,  even 
HCgHgOg.  Precipitation  is  hastened  by  rubbing  the 
side  of  the  vessel  with  a  glass  rod. 

2.  NH4OH  and  NaOH  precipitate  white  magnesium 
hydroxide,  Mg(0H)2,  soluble  in  acids  and  in  NH4CI. 

3.  Spectrum  (see  Table  VII,  p.  58). 

Ammonium  (salt  for  study,  ammonium  chloride,  NH4CI). 

1.  H2PtClg^  in  neutral  or  acid  solutions  precipitates 
yellow  crystalline  ammonium  platinic  chloride, 
(NH4)2PtClg,  insoluble  in  alcohol;  sparingly  soluble 
in  water  and  dilute  acids.  As  the  salt  is  somewhat 
soluble  in  water,  it  is  best  to  add  HgPtClg,  evaporate 
nearly  to  dryness,  and  add  a  few  drops  of  alcohol. 

2.  H2C4H40g  in  neutral  concentrated  solutions  precipi- 
tates white  acid  ammonium  tartrate,    H(NIl4)C4H40g, 

126 


METALS   OF  GROUP   VI  127 

insoluble  in  alcohol;  somewhat  soluble  in  water;  sol- 
uble in  acids.  The  reagent  should  be  added  like 
H^PtCV 

3.  NaOH  sets  free  NHg,  detected  by  the  odor,  with  red 
litmus  paper,  or  with  a  glass  rod  moistened  with  con- 
centrated HCl. 

4.  Heated  on  a  platinum  foil,  all  ammonium  salts  are 
completely  volatilized. 

Potassium  (salt  for  study,  potassium  chloride,  KCl). 

1.  HgPtClg  gives  a  reaction  very  similar  to  1,  under 
ammonium. 

2.  HgC^H^Og  gives  a  reaction  very  similar  to  2,  under 
ammonium. 

3.  Heated  in  a  non-luminous  flame,  potassium  salts 
color  the  flame  violet,  which  appears  purple  when  seen 
through  a  blue  glass. 

4.  Spectrum  (see  Table  VII,  p.  58). 

Sodium  (salt  for  study,  sodium  chloride,  NaCl). 

1.  H2K2Sb207  in  concentrated  neutral  solutions  precipi- 
tates sodium  pyro-antimoniate,  H2Na2Sb20y. 

2.  Heated  in  a  non-luminous  flame,  sodium  salts  color 
the  flame  yellow,  which  color  cannot  be  seen  through  a 
blue  glass. 

3.  Spectrum  (see  Table  VII,  p.  58). 

PROCESS   OF   SEPARATION 

As  the  group  is  composed  of  metals  not  bound  by  a 
common  group  reagent,  and  as  salts  of  some  of  the 


128  CHEMICAL  ANALYSIS 

metals  may  have  been  used  as  reagents  in  previous 
groups,  the  detection  of  the  individuals  cannot  be  made 
entirely  by  separations,  but  in  part  by  separations  and 
in  part  by  individual  tests  in  the  original  solutions. 

The  detection  of  the  members  of  this  group  is  based 
upon  the  facts  that  NH^MgPO^  is  insoluble  in  presence 
of  NH4OH  and  NH^Cl ;  that  K^ViQ\  and  NH^PtClg 
are  insoluble  in  alcohol  solutions ;  and  that  ammonium 
salts  are  dissociated  by  strong  alkalies,  and  completely 
volatilized  by  strong  heat. 

Concentrate  the  solution  by  evaporation.  Divide  the 
solution  into  two  unequal  parts. 

First  Part  (smaller).  —  Add  NH^OH,  NH^Cl,  and 
HNagPO^.  A  white  crystalline  precipitate  appearing 
immediately  or  after  some  time  confirms  NH^MgPO^. 

The  precipitate  should  be  crystalline,  and  if  the  sides  of  the 
vessel  are  scratched  with  a  glass  rod  the  crystals  will  adhere  to 
the  vessel  in  clusters  in  the  path  of  the  scratching.  As  the  crys- 
talline structure  of  the  precipitate  is  distinctly  characteristic  of 
NH^MgPO^,  it  is  often  desirable  to  confirm  this  with  the  micro- 
scope. A  simple  method  is  to  allow  the  crystals  to  grow  for  sev- 
eral hours.  Then  carefully  decant  off  the  liquid,  and  examine 
the  sides  of  the  glass  vessel  with  a  reading  glass  or  simple  micro- 
scope. If  the  precipitate  appears  flocculent  and  crystals  cannot  be 
detected,  examine  it  for»Al(OH)3  with  the  spectroscope.  A1(0H)3 
can  possibly  appear  at  this  stage  of  the  separation  because  it  is 
soluble  in  excess  of  NH^OH,  and  may  have  been  brought  over 
from  Group  III. 

Second  Part  (larger),  —  Evaporate  to  dryness  and 
expel  ammonium  salts  by  removing  the  mass  from 
the  dish,  and  heating  on  a  platinum  foil  till  no  white 
fumes  appear  immediately  after  removal  from  the  flame. 


METALS   OF  GROUP    VI  129 

Moisten  the  residue  with  HCl  and  test  for  potassium 
and  sodium  with  the  simple  flame  and  with  the  spectro- 
scope. Confirm  the  test  for  potassium  by  adding  to 
the  moistened  residue  some  H2PtClg  and  alcohol. 

Both  ammonium  and  potassium  salts  give  very  similar  precipi- 
tates with  HgPtClg  ;  hence  it  is  necessary  to  remove  all  the  ammo- 
nium salts  by  sublimation.  Care  should  be  taken  not  to  heat  the 
substance  too  much,  as  potassium  chloride  is  also  somewhat  vola- 
tile and  might  be  lost. 

As  ammonium  salts  are  used  as  reagents  in  several 
groups,  it  is  obviously  necessary  to  test  for  ammonia  in 
the  original  solution. 

Concentrate  some  of  the  original  solution  and  test 
with  NaOH  for  ammonium  salts. 


CHAPTER   XIII 

ACIDS  OF  GROUP  I:  CHROMIC,  CARBONIC,  SILICIC,  SUL- 
PHUROUS, SULPHURIC,  PHOSPHORIC,  BORIC,  OXALIC, 
TARTARIC,  AND  HYDROFLUORIC  ACIDS 

Characteristic  :  Insolubility  of  their  barium  salts  in  neutral 
solution. 

Group  Reagent  :  Barium  chloride. 

REACTIONS 

Chromic  Acid,  H2Cr04  (salt  for  study,  potassium  chro- 
mate,  K2Cr04). 

1.  BaClj  precipitates  yellow  barium  chromate,  BaCrO^, 
soluble  in  dilute  acids,  except  HgSO^. 

2.  HgS  reduces  CrOg  to  CrgOg ;  so  that  Cr  will  be 
detected  in  the  analysis  for  the  metals,  even  though  it 
originally  were  present  in  its  acid  state  of  oxidation. 

3.  Pb(C2H302)2  precipitates  yellow  lead  chromate, — 
chrome  yellow, — insoluble  in  HCgHgOg. 

(For  other  reactions,  see  Chromium  as  a  metal.) 

Carbonic  Acid,  H2CO3  (salt  for  study,  sodium  carbonate, 
Na2C03). 

1.  BaCl2  precipitates  white  barium  carbonate,  BaCOg, 
insoluble    in    water;   soluble    in    acids,  except  H2SO4.    > 
This  reaction  occurs  only  with  salts  of  HgCOg,  not  with 
the  free  acid, 

130 


ACIDS   OF  GROUP  I  131 

2.  HCl  and  other  acids,  excepting  HgS  and  HCN, 
decompose  carbonates  with  evolution  of  COg.  This 
gas  is  readily  soluble  in  water ;  and  in  dilute  solutions 
of  the  carbonates,  it  may  not  be  formed  in  sufficient 
quantity  to  oversaturate  the  solution  and  escape.  From 
concentrated  or  hot  solutions,  it  escapes  with  efferves- 
cence. Being  heavier  than  air,  it  may  be  detected  by 
decantation  into  a  test-tube  containing  lime  water,  its 
presence  being  shown  by  the  appearance  of  a  milky 
precipitate  :  — 

Ca(0H)2  +  CO2  =  CaCOg  +  H2O. 

An  excess  of  COg  will  dissolve  the  precipitate  first 
formed.  Sulphur  dioxide,  SO2,  will  also  produce  a 
white  precipitate  with  lime  water,  but  can  usually  be 
detected  by  its  odor. 

Silicic  Acid,   H4Si04  (salt    for    study,   sodium  silicate, 
Na4Si04). 

1.  BaCl2  precipitates  white  barium  silicate,  BagSiO^, 
decomposed  by  dilute  HCl,  with  separation  of  gelati- 
nous 1148104. 

2.  HCl  added,  drop  by  drop,  to  the  solution  of  a  sili- 
cate, precipitates  gelatinous  H^SiO^,  which,  on  evapora- 
tion to  dryness,  is  decomposed  with  the  formation  of 
silicic  anhydride,  SiOg. 

3.  Fused  with  NagCOg  on  a  platinum  foil  until  bub- 
bles of  gas  cease  to  escape,  most  insoluble  silicates  are 
changed  by  metathesis  to  sodium  silicate  and  a  metallic 
carbonate  or  oxide.  If  the  fused  mass  is  then  boiled 
with  dilute  HCl  and  filtered,  the  filtrate  will  contain 


132  CHEMICAL  ANALYSIS 

the  chloride  of  the  metal  and  the  residue  will  consist  of 
H^SiO^:  — 

(a)  BagSiO^  +  2Na2C03  =  2BaC03  +  Na^SiO^ ; 
(h)  BaCOg  +  2HC1  =  BaCla  +  H2O  +  COg; 
(e)  Na^SiO^  +  4HC1  =  H^SiO^  +  4NaCl. 

4.  HF  in  an  aqueous  solution,  or  in  gaseous  form, 
decomposes  SiOg  with  evolution  of  silicon  tetrafluoride, 
SiF^:  — 

SiOg  +  4HF  =  SiF^  +  2H2O. 

If  a  silicate  is  mixed  with  three  parts  of  NH^F  or 
five  parts  of  CaFg,  moistened  with  concentrated  HgSO^, 
and  then  heated  till  fumes  cease  to  escape,  the  silicic 
acid  is  decomposed  and  expelled :  — 

(a)  H2SO4  +  2NH4F  =  (NH4)2S04  +  2HF; 

(h)  6HF  +  Na2Si03     =  Na2SiFg  +  3  H2O ; 

(c)  Na2SiF6  +  H2S04  =  Na2S04  +  2HF  +  SiF4. 

6.  Metaphosphate  bead  dissolves  the  metallic  parts  of 
the  silicates,  but  not  the  Si02,  which  remains  floating 
in  the  fused  bead.  As  SiOg  is  not  affected,  the  outline 
of  the  particle  of  the  silicate  remains  intact,  giving 
rise  to  the  so-called  "skeleton  bead." 

ANALYSIS   OF   SILICATES 

There  are  two  classes  of  silicates  important  in  analytical  chem- 
istry —  silicates  decomposed  by  acids,  and  those  not  decomposed 
by  acids :  — 

First  class :  Silicates  decomposed  by  acids.  This  is  not  a  very 
numerous  class,  composed  for  the  most  part  of  the  soluble  alkali 
metal  silicates  and  a  few  less  soluble  single  and  double  silicates 
of  other  metals.     The  analysis  of  this  class  is  quite  simple.     This 


ACIDS    OF  GROUP  I  133 

is  accomplished  by  treatment  of  the  silicates  with  HCl,  which  by- 
metathesis  form  soluble  chlorides  of  the  metals  and  colloidal 
silicic  acid. 

Second  class  :  Silicates  not  decomposed  by  acids.  This  consti- 
tutes by  far  the  more  numerous  class,  including  the  natural  sili- 
cates. Many  natural  silicates  contain  the  alkali  metals  combined 
with  other  metals.  The  varieties  of  feldspar  are  representatives 
of  this  kind. 

The  analysis  of  silicates  riot  decomposed  by  acids  is  usually 
conducted  by  one  of  three  methods :  — 

Method  1.  —  Fusion  with  alkali-metal  carbonates.  By  metath- 
esis, soluble  silicates  and  carbonates  of  the  metals  of  the  original 
silicates  are  formed,  —  which  resulting  salts  are  then  decomposed 
by  HCI  (see  reactions  above).  Finely  powder  the  silicate,  mix 
with  about  3  parts  of  NagCOg,  or  fusion  mixture,  and  heat  to  quiet 
fusion  in  a  platinum  crucible  or  foil.  When  cool,  boil  the  mass 
in  water.  Filter,  evaporate  to  a  small  bulk,  and  add  concentrated 
HCl.  HgSiOg  will  precipitate  as  a  gelatinous  mass.  If  it  is  desired 
to  test  for  the  presence  of  alkali  metals  in  the  silicate,  this  method 
cannot  be  used,  as  the  carbonates  of  these  metals  are  added  as 
a  flux. 

Method  2} — (Method  of  J.  Lawrence  Smith.)  Fusion  with 
NH4CI  and  CaCOg.  An  insoluble  silicate  like  feldspar,  contain- 
ing alkali  metals,  may  be  converted  into  soluble  alkali-metal 
chlorides  and  some  insoluble  hydroxides,  by  heating  to  redness 
in  a  covered  platinum  crucible  with  1  part  NH^Cl  and  8  parts 
powdered  CaCOg.  In  all  fusions  it  is  necessary  for  both  the 
substance  and  the  flux  to  be  reduced  to  very  fine  powders,  and 
intimately  mixed. 

Method  3.  —  Fusion  with  BaO.  Fuse  in  a  platinum  crucible  a 
mixture  of  1  part  of  the  powdered  silicate  and  4  parts  BaO. 
Digest  the  mass  in  a  little  water  to  detach  it  from  the  crucible, 
and  then  dissolve  in  HCl.  Add  NH^OH  till  alkaline,  filter, 
evaporate  to  dryness,  and  ignite. 


134  CHEMICAL  ANALYSIS 

Sulphurous  Acid,  H2SO3  (salt  for  study,  sodium  sulphite, 

NaaSOs). 

1.  BaClj  precipitates  white  barium  sulphite,  BaSOg, 
soluble  in  dilute  HCl. 

2.  Nascent  hydrogen  reduces  sulphites  to  sulphides, 
which  are  decomposed  by  an  excess  of  HCl  with  evolu- 
tion of  HgS,  detected  by  its  odor,  or  with  Pb{C2H302)2- 

3.  HgS  decomposes  sulphites  with  separation  of 
sulphur. 

4.  HCl  decomposes  sulphites  with  evolution  of  SOg, 
detected  by  its  odor  and  by  the  production  of  a  white 
precipitate  of  calcium  sulphite  with  lime  water. 

Sulphuric  Acid,  H2SO4  (salt  for  study,  sodium  sulphate, 
Na2S04). 

1.  BaCl2  precipitates  white  barium  sulphate,  BaSO^, 
insoluble  in  water  or  acids ;  decomposed  by  fusion  with 
NagCOg  in  a  platinum  crucible  or  foil,  sodium  sulphate 
and  barium  carbonate  being  formed :  — 

NaaCOg  +  BaSO^  =  Na2S04  +  BaCOg. 

In  like  manner,  the  other  insoluble  sulphates,  SrS04, 
CaSO^,  and  PbS04,  are  decomposed  by  fusion  with 
Na2COg  or  by  boiling  with  its  solution. 

2.  Pb(C2Hg02)2  precipitates  white  lead  sulphate,  PbS04, 
almost  insoluble  in  dilute  HNOg ;  soluble  in  hot  con- 
centrated HCL' 

3.  Fused  with  NagCOg  on  charcoal,  sulphates  are 
reduced  to  sulphides.  If  the  mass  is  moistened  with 
very  dilute  HCl  and  placed  on  a  bright  silver  coin,  the 
latter  will  be  stained  black. 


ACIDS    OF  GROUP  I  135 

Phosphoric  Acid,  H3PO4  (salt  for  analysis,  sodium  phos- 
phate, HNagPOi). 

1.  BaCl2  precipitates  white  barium  phosphate,  HBaPO^, 
—  or  Ba3(P04)2,  if  the  solution  contained  a  normal  phos- 
phate, —  soluble  in  HCl  and  HNO3. 

2.  MgSO^  in  presence  of  NH^OH  and  NH^Cl  precipi- 
tates white  crystalline  ammonium  magnesium  phos- 
phate, NH^MgPO^,  soluble  in  acids.  (Compare  with 
NH^MgAsO^.) 

3.  (NH4)2Mo04  in  HNO3  solution  precipitates,  in 
the  cold,  yellow  ammonium  phospho-molybdate, 
(Mo03)i2*(-^H4)3P04.  (Compare  with  behavior  of  the 
same  reagent  toward  arsenates.) 

4.  FeCl3  in  presence  of  NaC2H302  precipitates  yellow 
ferric  phosphate,  FePO^,  soluble  in  strong  acids  and 
excess  of  FeCl3;  insoluble  in  HC2H3O2. 

Boric  Acid,  H3BO3  (salt  for  study,  borax,  Na2B407). 

1.  BaCl2  precipitates  white  sodium  barium  borate, 
Na2Ba5(B03)4,  soluble  in  acids,  except  H2SO4. 

2.  H2SO4  precipitates  from  hot  solutions  of  borates,  on 
cooling,  crystalline  boric  acid,  H3BO3. 

3.  Alcohol,  added  to  free  boric  acid  or  to  a  borate  with 
concentrated  HgSO^  and  then  kindled,  burns  with  a 
green  flame,  especially  upon  stirring  the  mixture. 

4.  Turmeric  paper,  immersed  in  a  slightly  acid  (HCl) 
solution  of  boric  acid  or  a  borate  and  then  dried, 
shows  a  reddish  tint  which  is  turned  blue  by  NaOH. 


136  CHEMICAL  ANALYSIS 

Oxalic  Acid,   H2C2O4  (salt   for  study,   sodium  oxalate, 
Na2C204). 

1.  BaCl2  precipitates  from  neutral  solutions  white 
barium  oxalate,  BaCgO^,  somewhat  soluble  in  dilute 
NH4CI  and  many  organic  acids;  soluble  in  HCl  and 
HNO3. 

2.  Lime  water  and  soluble  calcium  salts  precipitate  white 
calcium  oxalate  CaCgO^,  soluble  in  HCl  and  HNO3; 
insoluble  in  organic  acids. 

3.  Concentrated  H^SO^,  heated  with  oxalic  acid  or  an 
oxalate,  removes  water,  and  the  compound  is  decom- 
posed into  CO2  and  CO  :  — 

H2C2O4  H-  H2SO4  =  CO2  +  CO  +  H2S04-H20. 

If  in  sufficient  quantity  the  CO  gas  can  be  burned  with 
its  characteristic  blue  flame. 

4.  Heating  decomposes  all  oxalates  with  formation  of 
carbonates  or  oxides  of  the  metals,  and  evolution  of  CO 
or  CO2. 

Tartaric  Acid,  H2C4H4O6  (salt  for  study,  potassium  tar- 
trate, K2C4H4O6). 

1.  BaClg  (or,  better,  CaClg)  from  neutral  solutions 
precipitates  white  barium  (or  calcium)  tartrate,  soluble 
in  acids,  except  H2SO4. 

2.  AgNOg  precipitates  white  silver  tartrate, 
Agfi^fif^^  soluble  in  NH^OH.  On  warming  this 
solution,  black  metallic  silver  is  deposited.  If  the 
Ag2C4H40g  be  carefully  redissolved  in  the  least  possible 
amount  of  NH^OH,  and  if  this  solution  be  heated  gently 


ACIDS   OF  GROUP  I  137 

in  a  test-tube,  a  mirror  of  metallic  silver  will  be  depos- 
ited on  the  walls  of  the  tube.  AgNOg  precipitates 
AggC^H^Og  only  from  neutral  solutions.  This  reaction 
distinguishes  tartaric  from  most  other  organic  acids. 

3.  Heated  in  a  closed  tube,  tartrates  char  and  emit 
inflammable  vapors  with  the  odor  of  burnt  sugar. 
Commingled  with  the  carbon  residue  is  also  a  carbon- 
ate, detected  by  effervescence  on  adding  HCl. 

Hydrofluoric  Acid,  HF  (salt  for  study,  ammonium  fluoride, 

NH4F). 

1.  BaCl2  precipitates  white  barium  fluoride,  BaFg, 
soluble  with  difliculty  in  HCl  and  HNO3. 

2.  Concentrated  H^SO^  mixed  to  a  paste  with  powdered 
fluorides  and  warmed  in  a  platinum  vessel  expels  gas- 
eous HF :  — 

2NH4F  +  H2SO4  =  (NH4)2S04  +  2HF. 

If  the  vessel  is  loosely  covered  for  an  hour  with  a 
watch-glass  which  previously  has  been  coated  with  wax 
through  which  some  lines  have  been  cut  with  a  sharp 
instrument,  the  lines  will  be  seen  to  have  been  etched 
into  the  glass  upon  removal  of  the  wax.  The  reaction 
involved  is  identical  with  No.  4,  under  silicic  acid. 

DETECTION  OF  THE   ACIDS   OF   GROUP  I 

The  analysis  for  acids  cannot  be  made  by  following  a 
systematic  scheme  of  separation,  such  as  is  used  in  the 
analysis  for  metals ;  on  the  contrary,  the  presence  or 
absence  of  each  acid  must  be  established  chiefly  by 
individual  tests  applied  to  the  original  material. 


138  CHEMICAL  ANALYSIS 

For  convenience  the  members  of  Group  I  may  be 
classified  as  follows  :  — 

Sub-group  J_  H2Cr04,  HgCOg,  H^SiO^,  H2SO3. 
These  acids  are  decomposed,  in  solution,  by  HCl  and 

Sub-group  ZT— H2SO4,  H3PO4,  H3BO3,  H2C2O4, 
H^C^H^Oq^  HF.  These  acids  are  not  decomposed  by 
HCl  or  H2S. 

Neutralize  a  small  portion  of  the  original  solution, 
and  add  some  BaClg  (or  Ba(N03)2,  if  metals  of  Group  I 
are  present).  A  precipitate  confirms  the  presence  of 
one  or  more  acids  of  Group  I.  Divide  a  larger  portion 
of  the  original  solution  into  four  parts :  — 

Part  I^  for  chromic  add.  —  A  yellow  color  indicates 
chromic  acid,  confirmed  by  acidifying  with  HC2H3O2 
and  adding  Pb(C2H302)2- 

Part  II,  for  carbonic  acid,  —  Add  HCl  and  warm. 
An  effervescence  of  an  odorless  gas  indicates  the 
presence  of  CO2.  Confirm  by  testing  with  lime 
water. 

Part  in,  for  sulphurous  acid.  —  Add  HCl  and  warm. 
Effervescence  with  odor  of  burning  sulphur  indicates 
the  presence  of  SO2,  confirmed  by  passing  the  gas 
through  lime  water. 

Part  IV,  for  silicic  acid.  —  Add  dilute  HCl,  drop  by 
drop.  A  gelatinous  precipitate  indicates  H^SiO^,  con- 
firmed by  evaporating  to  dryness  and  testing  with  the 
metaphosphate  bead. 

If  any  of  these  acids  are  present,  they  must  be 
removed  from  solution  before  testing  for  the  members 
of  Sub-group    II.     H2Cr04   is   destroyed   by    HgS,   in 


ACIDS    OF  GROUP  I  139 

presence  of  HCl ;  HgCOg  and  H2SO3  are  driven  off  by 
boiling  with  HCl;  and  1148104  is  removed  by  evapo- 
ration with  HCl. 

The  solution,  thus  freed  of  members  of  Sub-group  I, 
is  now  neutralized  exactly  with  NH4OH,  —  free  of 
(NH4)2C03,  —  and  its  examination  is  continued  as 
follows  :  — 

To  a  small  portion  BaClg  is  added.  If  no  precipitate 
is  formed,  all  members  of  Sub-group  II  are  absent.  If 
a  precipitate  is  formed  which  dissolves  on  the  addition 
of  HCl,  H2SO4  is  absent,  but  other  members  may  be 
present.  If  a  precipitate  is  formed  which  does  not 
dissolve  in  HCl,  HgSO^  (possibly  other  acids)  is  present. 
In  either  of  the  latter  cases  it  is  necessary  to  test  indi- 
vidually for  the  remaining  acids  of  the  group  in  small 
portions  of  the  solution. 

Part  I,  for  phosphoric  acid, — Add  a  few  drops  of  the 
solution  to  a  strong  HNO3  solution  of  (NH4)2Mo04,  and 
warm  gently.  A  yellow  crystalline  precipitate  confirms 
the  presence  of  H3PO4. 

Part  II,  for  boric  acid. — Acidify  some  of  the  solution 
with  HCl  and  test  with  turmeric  paper.  Evaporate 
another  portion  almost  to  dryness,  add  alcohol  and  con- 
centrated H2SO4,  and  kindle.  A  green  flame  confirms 
the  presence  of  H3BO3. 

Part  III,  for  oxalic  acid. — Add  lime  water  and  boil 
the  white  precipitate  with  HC2H3O2.  If  the  precipitate 
fails  to  dissolve,  it  confirms  the  presence  of  H2C2O4. 

Part  IV,  for  tartaric  acid.  —  Neutralize  the  solution 
and  add  CaCl2.  If  a  white  precipitate  occurs,  filter,  dry 
the  residue,  and  heat  in  a  closed  tube.     Charring  with 


140  CHEMICAL  ANALYSIS 

the  odor  of  burnt  sugar,  and  effervescence  of  the  residue 
with  HCl,  confirm  the  presence  of  HgC^H^Og. 

Part  Vy  for  hydrofluoric  acid. — Evaporate  the  solution 
to  dryness,  transfer  the  residue  to  a  platinum  crucible, 
add  concentrated  HgSO^,  and  cover  with  a  watch-glass. 
If  the  gas  etches  the  glass  cover,  the  presence  of  HF 
is  confirmed. 


CHAPTER   XIV 

ACIDS  OF  GROUP  II:  HYDROCHLORIC,  HYDROBROMIC, 
HYDRIODIC,  HYDROCYANIC,  HYDROFERROCYANIC, 
HYDROFERRICYANIC,  SULPHOCYANIC,  AND  HYDRO- 
SULPHURIC    ACIDS 

Characteristic  :  Insolubility  of  their  silver  salts  in  dilute 
nitric  acid. 

Group  Reagent  :  Silver  nitrate. 

REACTIONS 

Hydrochloric  Acid,  HCl  (salt  for  study,  sodium  chloride, 

NaCl). 

1.  AgNOg  precipitates  white  silver  chloride,  AgCl, 
insoluble  in  dilute  acids;  soluble  in  KCN,  NH^OH, 
and  in  boiling  solution  of  ammonium  "sesqui"  car- 
bonate. 

2.  Pb02  or  Mn02  with  concentrated  H2SO4  expels  chlo- 
rine gas,  detected  with  starch-KI  paper. 

3.  K2Cr207  with  concentrated  H2SO4  gives  red 
fumes,  condensing  to  a  brown  liquid,  chromic  oxy- 
chloride,  Cr02Cl2,  changing  to  yellow  (NH4)2Cr04 
on  the  addition  of  NH^OH.  The  dry  chloride 
should  be  triturated  with  K2Cr207  crystals,  and  dis- 
tilled with  concentrated  HgSO^  in  a  small  retort 
(25  c.c). 

141 


142  CHEMICAL  ANALYSIS 

Hydrobromic  Acid,  HBr  (salt  for  study,  potassium 
bromide,  KBr). 

1.  AgNOg  precipitates  yellow  silver  bromide,  AgBr, 
insoluble  in  dilute  acids  and  in  ammonium  "sesqui" 
carbonate;  soluble  in  NH^OH  and  KCN. 

2.  PbOg  with  concentrated  HgSO^  expels  brown  vapors 
of  bromine,  identified  by  their  odor  and  color. 

3.  'Kji^rjdrj  with  concentrated  H2SO4  expels  bromine, 
which  is  decolorized  by  NH^OH,  forming  NH^Br. 

4.  Chlorine  liberates  bromine,  detected  in  small  quan- 
tities by  coloring  carbon  disulphide  or  chloroform 
brownish-red.  Mix  the  bromide  solution  with  about 
1  c.c.  of  CSg,  then  add  dilute  chlorine  water,  drop  by 
drop,  and  shake  well.  The  globules  of  CSg  will  assume 
a  reddish  tint.  An  excess  of  chlorine  should  be  avoided, 
lest  it  combine  with  bromine  to  form  colorless  bromine 
chloride,  BrCl. 

Hydriodic  Acid,  HI  (salt  for  study,  potassium  iodide,  KI). 

1.  AgNOg  precipitates  yellow  silver  iodide,  Agl, 
insoluble  in  dilute  acids,  NH^OH  and  ammonium 
"sesqui"  carbonate;  soluble  in  KCN. 

2.  PbOg  with  concentrated  HC2H3O2  liberates  violet 
iodine,  turning  starch  paper  blue. 

3.  KgCrgO^  with  concentrated  H2SO4  liberates  iodine. 

4.  Chlorine  water  liberates  iodine,  turning  starch 
paper  blue.  An  excess  of  chlorine  will  decolorize 
the  paper  by  formation  of  iodine  chloride,  ICl. 

5.  KNO2  in  concentrated  H2SO4  liberates  iodine. 
Into  a  clear  solution  of  starch  paste  and  an  iodide,  dip 


ACIDS   OF  GBOUP  II  143 

a  glass  rod  moistened  with  a  solution  of  KNO2  in  con- 
centrated H2SO4.  The  liquid  in  contact  with  the  rod 
becomes  blue.  It  is  necessary  to  keep  the  reagent  cold, 
as  iodized  starch  becomes  colorless  in  hot  water. 

Hydrocyanic  Acid,  HCN  (salt  for  study,  potassium  cya- 
nide, KCN). 

1.  AgNOg  precipitates  white  silver  cyanide,  AgCN, 
soluble  in  excess  of  KCN,  forming  the  salt  KAg(CN)2. 
AgCN  is  also  soluble  in  NH^OH  and  boiling  HCl. 

2.  FeSO^,  with  a  few  drops  of  FeClg,  added  to  tlie  solu- 
tion of  a  cyanide  in  weak  NaOH,  precipitates  a  bluish- 
green  mixture  of  ferrous  ferric  hydroxide,  Fe302(OH)^, 
and  Prussian  blue.  Fe302(OH)4  can  be  dissolved  with 
dilute  HCl,  leaving  the  Prussian  blue  intact. 

3.  (NH4)2S3,(a  few  drops)  and  a  drop  of  NaOH  added 
to  a  cyanide  solution,  form  ammonium  sulphocyanate, 
NH^CNS,  on  heating.  Evaporate  the  solution  to  dry- 
ness and  test  by  dissolving  in  dilute  HCl  and  adding 
FeClg  solution.  A  deep  red  coloration  shows  the  pres- 
ence of  HCNS,  derived  from  HCN  :  — 

(NH4)2S^  +  4  KCN  =  4  KCNS  +  (NH4)2S(rc  -  4). 

4.  HNaCOg  heated  with  a  cyanide  expels  HCN  gas, 
identified  by  its  odor  and  the  rose  color  of  its  flame. 

Hydroferrocyanic  Acid,  H4Fe(CN)Q  (salt  for  study,  potas- 
sium ferrocyanide,  K4Fe(CN)e). 

1.  AgNOg  precipitates  white  silver  ferrocyanide, 
Ag4Fe(CN)g,  soluble  in  KCN;  insoluble  in  NH^OH 
and  HNOg. 


144  CHEMICAL  ANALYSIS 

2.  FeClg  precipitates  Prussian  blue  (see  Iron,  p.  107). 

3.  CuSO^  precipitates  brown  cupric  ferrocyanide, 
Cu2Fe(CN)6  (see  Copper,  p.  84). 

Hydroferricyanic  Acid,  H3Fe(CN)6  (salt  for  study,  potas- 
sium ferricyanide,  K3Fe(CN)g). 

1.  AgNOg  precipitates  orange-red  silver  ferricyanide, 
Ag3Fe(CN)6,  soluble  in  NH^OH  and  KCN;  insoluble 
in  HNO3. 

2.  FeSO^  precipitates  TurnbuU's  blue  (see  Iron,  p.  107). 

Sulphocyanic  Acid,  HCNS  (salt  for  study,  potassium 
sulphocyanate,  KCNS). 

1.  AgN03  precipitates  white  silver  sulphocyanide, 
AgCNS,  soluble  in  NH^OH;  insoluble  in  dilute 
HNO3. 

2.  FeClg  acidified  with  HCl  gives  a  deep  red  colora- 
tion of  Fe(CNS)3  (see  Iron,  p.  108). 

Hydrosulphuric  Acid,  H2S  (salt  for  study,  sodium  sul- 
phide, NagS). 

1.  AgN03  precipitates  black  silver  sulphide,  Ag^S. 

2.  Na2FeNO(CN)5  (sodium  nitro-prusside)  added  to 
alkaline  (NaOH)  solution  of  a  sulphide  gives  a  brilliant 
red-violet  tint. 

3.  Fused  with  NaOH,  insoluble  sulphides  form  NagS  ; 
and  on  dissolving  the  mass  in  a  little  water,  the  solu- 
tion will  tarnish  a  bright  silver  coin  brown. 

4.  HCl  sets  free  H2S  from  all  soluble,  and  from  many 
insoluble  sulphides ;  recognized  by  its  odor  and  by  its 


1 


ACIDS   OF  GROUP  II  145 

power  of  blackening  paper  moistened  with  a  solution  of 

DETECTION  OF  THE   ACIDS   OF  GROUP  II 

The  separation  and  identification  of  the  acids  of  this 
group  are  accomplished  by  the  following  means :  — 

(a)  The  removal  of  HgS  by  means  of  a  solution  of 
ZnSO^  in  NaOH. 

(b)  Hager's  method  of  detecting  HCl,  HBr,  and  HI 
in  the  presence  of  each  other ;  based  upon  the  different 
degrees  of  solubility  of  AgCl,  AgBr,  and  Agl  in  ammo- 
nium "sesqui"  carbonate  and  NH^OH. 

(c)  The  detection  of  HON  in  the  absence  of 
H4Fe(CN)6,  H3Fe(CN)6,  and  HCNS,  by  the  precipita- 
tion of  Prussian  blue  from  a  solution  of  a  cyanide  by 
FeSO^,  FeClg,  and  NaOH. 

(d)  The  detection  of  HON  in  the  presence  of 
H4Fe(CN)6,  H3Fe(CN)6,  and  HCNS,  by  the  evolution 
of  HCN  on  distilling  with  HNaCOg. 

HgS  must  first  be  tested  for  in  a  small  portion  of  the 
original  solution,  preferably  by  adding  HCl,  boiling, 
and  noting  whether  any  gas  is  given  off  which  causes 
lead  acetate  paper  to  blacken.  If  found,  it  must  be 
removed  from  the  remainder  of  the  solution  before  test- 
ing for  the  other  members  of  the  group,  since  its  pres- 
ence would  hinder  their  detection.  Therefore,  treat  a 
sufficient  portion  of  the  solution  with  a  solution  of 
ZnSO^  in  an  excess  of  NaOH,  which  will  precipitate 
the  H2S  as  ZnS.  Reject  the  precipitate,  and  divide  the 
filtrate,  or  portion  of  the  original  solution  if  HgS  is 
absent,  into  three  parts. 


146  CHEMICAL  ANALYSIS 

Part  I,  for  HCl,  HBr,  and  iTZ— Acidify  with  HNO3 
and  add  AgNOg.  Filter  and  reject  the  filtrate.  Boil 
the  residue  with  100  parts  of  a  solution  of  ammonium 
"  sesqui "  carbonate.  Decant  the  clear  supernatant 
liquid,  add  more  ammonium  "  sesqui "  carbonate,  and 
again  boil  and  decant.  The  decanted  liquid  may  con- 
tain AgCl,  which  can  be  determined  by  acidifying  with 
HNO3.  The  residue  from  which  the  liquid  has  been 
decanted  may  consist  of  AgBr  and  Agl.  Treat  it  with 
a  dilute  solution  of  NH^OH  (5  per  cent  ammonia  water) 
and  filter.  The  filtrate  may  contain  AgBr,  detected 
by  acidifying  with  HNO3.  "^^^  residue  may  be  Agl, 
indicated  by  its  yellow  color.  For  a  further  confirma- 
tion of  the  three  halogens  (consisting  of  the  AgCl  and 
AgBr  precipitates  from  the  ammoniacal  solutions  and 
the  undissolved  Agl)  each  can  be  fused  with  NagCOg, 
boiled  with  water,  and  filtered :  — 

2 AgCl  +  Na2C03  =  2NaCl  +  AggCOg,  etc. 

The  filtrates  can  be  tested  for  the  individual  halogens 
as  follows :  — 

(a)  Solution  of  NaCl.  Evaporate  to  dryness  and  heat 
with  concentrated  HgSO^  and  PbOg.  The  evolved  chlo- 
rine can  be  detected  by  its  odor,  its  bleaching  moistened 
litmus  paper,  or  its  effect  on  starch-KI  paper. 

(b)  Solution  of  NaBr.  Evaporate  to  dryness  and 
heat  with  concentrated  H2SO4  and  PbOg.  The  evolved 
bromine  can  be  detected  by  its  odor  or  by  its  color. 

{c)  Solution  of  Nal.  Neutralize  with  HNO3  and  add 
some  drops  of  starch  paste  and  chlorine  water.  A  blue 
solution  confirms  presence  of  the  iodide. 


i 


ACIDS   OF  GROUP  II  147 

Part  II  for  ff^FeiCN),,  HsFe^CN)^,  and  EONS,— 
Neutralize  with  HNO3  and  divide  into  two  small  parts. 
Pour  one  part  into  a  test-tube  and  shake  the  tube  so 
that  its  sides  will  be  moistened  with  the  liquid.  Hold- 
ing the  tube  obliquely,  add  a  few  drops  of  dilute  FeClg 
solution  so  that  they  will  run  down  the  sides  of  the 
tube.  A  red  coloration  indicates  the  presence  of  HCNS. 
If  H4Fe(CN)g  is  present,  Prussian  blue  will  be  formed 
also,  but  the  red  color  can  be  seen  commingled  with  the 
blue.  Add  more  FeClg.  The  formation  of  Prussian 
blue  confirms  presence  of  H4Fe(CN)g.  To  the  second 
smaller  part  add  FeSO^.  The  formation  of  Turnbull's 
blue  confirms  the  presence  of  H3Fe(CN)g. 

Part  III,  for  HCN.  —  li  }1^Yq{C^)^,  H3Fe(CN)g, 
and  HCNS  are  absent,  add  NaOH,  FeSO^,  a  few  drops 
of  FeClg,  and  HCl  in  excess.  Formation  of  Prussian 
blue  confirms  the  presence  of  HCN. 

If  H4Fe(CN)6,  H3Fe(CN)6,  and  HCNS  are  present, 
add  some  solid  bicarbonate  of  sodium,  HNaCOg,  to  the 
neutral  solution  in  a  test-tube,  and  boil.  The  odor  of 
bitter  almonds  indicates  the  presence  of  HCN.  Con- 
firm by  kindling  the  gas.  It  should  burn  with  a  rose- 
tinted  flame. 

HCN  is  a  deadly  poison  ;  do  not  inhale. 


CHAPTER   XV 

ACIDS  OF  GROUP  HI :  NITRIC,  CHLORIC,  AND  ACETIC  ACIDS 
No  group  characteristic.    No  group  reagent. 

REACTIONS 

Nitric  Acid,  HNO3  (salt  for  study,  potassium  nitrate, 
KNO3). 

1.  Heated  on  charcoal,  nitrates  deflagrate  with  igni- 
tion, giving  off  COg  :  — 

2KNO3  +  C  =  2KNO2  +  CO2. 

Use  small  quantities  of  the  nitrate  in  performing  this  experiment. 

2.  Heated  with  KCN  in  a  platinum  crucible  or  foil, 
nitrates  deflagrate  with  ignition  and  detonation :  — 

KNO3  +  KCN  =  KNO2  +  KCNO. 

3.  Mixed  with  copper  filings  and  heated  with  con- 
centrated H2SO4,  nitrates  give  red  fumes  of  NO2. 

4.  If  a  concentrated  solution  of  FeSO^,  free  of  ferric 
salts,  be  carefully  added  to  the  cold  solution  of  a 
nitrate  in  concentrated  H2SO4,  so  that  the  two  solu- 
tions form  distinct  layers,  a  brown  ring  will  be  formed 
at  their  junction,  (FeS04)2NO  :  — 

(a)  2HN03  +  6FeS04  +  3H2S04 

=  3Fe2(S04)3  +  4H20  +  2N0; 

(b)  2FeS04  +  NO  =  (FeS04)2NO. 


ACIDS   OF  GROUP  III  149 

5.  Brucine  dissolved  in  concentrated  H2SO4  gives  a 
deep  red  color  with  nitrates.  Touch  the  edge  of  the 
dissolved  brucine  with  a  glass  rod  moistened  with 
nitrate  solution ;  a  distinct  red  ring  will  bound  the  rod. 

6.  Reduced  with  zinc  dust  and  HgSO^,  nitrates  yield 
nitrous  acid,  HNO2,  detected  by  starch-KI  paper. 

7.  NaOH  with  zinc  dust  and  iron  filings,  on  heating, 
reduces  nitrates  and  sets  NHg  free  :  — 

HNO3  +  8  H  =  NH3  +  3  H2O. 

Chloric  :At:id,  HCIO3  (salt  for  study,  potassium  chlorate, 

KCIO3). 

1.  Heated  on  charcoal,  chlorates  deflagrate  with  vivid 
ignition,  giving  off  COg  :  — 

2KCIO3  +  3C  =  2KC1  +  3CO2. 

2.  Heated  with  KCN  in  a  platinum  crucible,  chlorates 
deflagrate  with  ignition  and  detonation  : —     . 

KCIO3  4-  3  KCN  =  KCl  +  3KCN0. 

As  HCIO3  gives  up  more  oxygen  than  HNO3,  the  chemical  action 
in  Reactions  1  and  2  is  necessarily  more  vigorous  than  in  those 
under  HNO3.     Therefore,  use  very  small  quantities  of  chlorate. 

3.  Concentrated  H2SO4  (a  few  drops),  added  with  a 
pipette  to  a  watch-glass  containing  a  chlorate  solution, 
liberates  chlorine  peroxide :  — 

3  KCIO3  +  2  H2SO4  =  KCIO4  +  2  CIO2  +  H2O  +  2HKSO4. 

The  peroxide  is  characterized  by  a  disagreeable  odor  and  a 
yellow  coloration ;  also  by  bleaching  a  blue  solution  of  indigo. 
Neither  heat  nor  large  quantities  of  reagents  should  be  used. 

4.  Brucine  behaves  very  much  alike  towards  nitrates 
and  chlorates. 


150  CHEMICAL  ANALYSIS 

Acetic  Acid,  HC2H3O2  (salt  for  study,  sodium  acetate, 
•       NaC2H302). 

1.  Heated  to  redness,  acetates  are  decomposed  with 
the  formation  of  carbonates  and  of  acetone,  CgHgO,  a 
liquid  of  penetrating,  pleasant,  ethereal  odor :  — 

2NaC2H302  =  Na2C03  +  CgHgO. 

2.  FeClg,  a  few  drops,  added  to  a  neutral  acetate 
solution,  produces  a  deep  red  coloration,  due  to  the 
formation  of  ferric  acetate,  Fe(C2H302)3.  On  boiling, 
the  solution  is  decolorized,  and  brown  basic  ferric  acetate 
is  precipitated. 

3.  Heated  with  concentrated  H2SO4  and  alcohol,  ace- 
tates yield  ethyl  acetate,  (C2H5)C2H302,  characterized 
by  its  pungent  ethereal  odor :  — 

NaC2H302  +  C2H5OH  =  (C2H5)C2H302  +  NaOH. 

DETECTION  OF   THE   ACIDS   OF   GROUP  III 

Individual  tests  must  be  made  for  the  three  acids  of 
this  group,  in  separate  portions  of  the  original  solution. 

If  iodides  or  bromides  are  present,  they  must  be 
removed  from  the  portion  which  is  to  be  tested  for 
HNO3  by  adding  HgClg  solution  and  filtering,  rejecting 
the  precipitate.  Otherwise,  they  would  give  a  dark 
coloration  on  the  addition  of  H2SO4.  In  testing  for 
HNO3,  Reactions  3  and  4  are  to  be  used. 

In  testing  for  KCIO3,  Reaction  3  is  to  be  employed. 
If  H3PO4  is  present,  it  is  to  be  removed  before  testing 
for  HC2H3O2,  since  it  forms  insoluble  FePO^  with 
FeClg.  (See  p.  111.)  Use  Reaction  2  in  testing  for 
HOgHgOg. 


CHAPTER  XVI 

THE  SYSTEMATIC  PROCEDURE  OF  ANALYSIS 
PRELIMINARY  TESTS 

The  physical  properties  of  the  substance  under  exami- 
nation— color,  odor,  whether  solid  or  liquid,  etc. — are 
first  to  be  noted. 

Solids.i  —  If  the  substance  is  a  solid,  apply  the  follow- 
ing tests  to  small  portions :  — 

(a)  Blowpipe  flame  on  charcoal  (see  Tables  IV  and 

V,  pp.  48,  49). 

(b)  Heating  in  a  closed  tube  (see  Table  II,  p.  45). 

(c)  Fusion  with  borax  bead  (see  Table  III,  p.  47). 

(d)  Flame  coloration  on  a  platinum  wire  (see  Table 

VI,  p.  51). 

(e)  Spectra  (see  Table  VII,  p.  58). 

A  larger  portion  of  the  solid  is  to  be  used  for  solu- 
tion, in  preparation  for  the  analysis  by  the  wet  way. 
First,  treat  it  with  water,  determining  whether  the 
whole  or  only  a  part  dissolves.  If  it  be  insoluble,  or 
only  partly  soluble,  divide  the  mixture  into  three  parts, 
two  small  and  one  large,  which  may  be  numbered 
respectively  1,  2,  and  3. 

To  1,  in  a  test-tube,  add  some  dilute  HCl  and  boil. 
If  still  insoluble  or  partly  insoluble,  add  an  equal  volume 
of  concentrated  HCl  and  boil  again.  If  soluble,  then 
treat  3,  the  largest  portion,  with  concentrated  HCl  and 

151 


152  CHEMICAL  ANALYSIS 

boil.  If  insoluble  or  partly  insoluble,  treat  2,  first  with 
dilute,  then  with  concentrated  HNO3.  If  soluble,  treat 
3  in  like  manner. 

If  insoluble  or  partly  insoluble  in  HNO3  as  well  as  in 
HCl  and  water,  combine  the  strong  HCl  and  HNO3 
mixtures,  1  and  2,  and  boil.  If  soluble,  treat  3  with 
aqua  regia.  If  insoluble  or  partly  insoluble,  recall 
which  of  the  four  solvents  —  water,  HCl,  HNOg,  or 
aqua  regia  —  dissolved  the  substance  most ;  and  treat 
the  larger  portion,  3,  with  that  solvent.  Filter.  Fuse 
the  residue  ^  with  fusion  mixture  on  a  platinum  foil 
or  in  a  platinum  crucible,  and  boil  the  mass  with 
water.  Sometimes  this  solution  can  be  added  directly 
to  the  filtrate  without  precipitation.  Generally,  how- 
ever, a  precipitate  will  be  formed.  In  order  to  deter- 
mine this,  take  small  portions  of  both  liquids  and  mix 
them.  If  no  precipitate  forms,  combine  the  whole  of 
both  solutions.  If  a  precipitate  forms,  separate  analyses 
must  be  made  of  the  two  solutions.^ 

Liquids  are  to  be  tested  with  litmus  paper  to  determine 
whether  they  are  neutral,  acid,  or  alkaline;  and,  also, 
small  portions  are  to  be  evaporated  to  dryness  on  the 
wjiter  bath.  No  residue  being  left,  a  neutral  reaction 
indicates  that  only  water  is  present;  whereas  an  acid 
or  alkaline  reaction  indicates  the  presence  of  a  volatile 
acid  or  of  ammonia. 

If  a  residue  is  left  on  evaporation:  — 

(a)  A  neutral  reaction  indicates  the  presence  in  solu- 
tion of  a  neutral  salt. 

(b)  An  acid  solution  may  be  either  (1)  a  free  acid, 
(2)   an   aqueous  solution   of   certain   normal  salts   likej 


SYSTEMATIC  PROCEDURE   OF  ANALYSIS       153 

FeClg  or  CuSO^,  which  have  acid  reactions,  (3)  certain 
acid  salts  like  bisulphate  of  potassium,  HKSO^,  or  (4) 
an  acid  solution  of  certain  salts. 

(c)  An  alkaline  solution  may  contain  (1)  a  free  alkali, 
(2)  certain  normal  salts  like  Na2C03  which  have  an  alka- 
line reaction,  or  (3)  an  alkaline  solution  of  certain  salts. 
In  theory  an  aqueous  solution  of  a  basic  salt  should 
react  alkaline ;  but  as  the  metals  which  form  basic  salts 
have  not  a  very  pronounced  metallic  character,  their 
alkalinity  is  either  too  weak  to  be  detected  by  litmus 
or,  being  very  weak,  is  neutralized  by  water. 

SYSTEMATIC   ANALYSIS   FOR   METALS 

If  the  solution  1  is  neutral  or  alkaline,^  add  dilute  HCl 
till  acid ;  if  acid,  boil  off  the  excess  of  acid,  and  when 
cold  add  dilute  HCl.  If  a  precipitate  is  formed,  filter, 
and  wash  the  residue  with  cold  water.  Analyze  the 
residue  for  members  of  Group  I,  according  to  the  direc- 
tions on  p.  80. 

Acidify  the  filtrate  strongly  with  more  HCl,  warm  to 
about  70°,  and  pass  a  constant  stream  of  HgS  through  it 
for  about  fifteen  minutes.  Then  cool  and  dilute  the  solu- 
tion, and,  before  filtering,  pass  HgS  again  till  saturation 
is  completed.  If  a  precipitate  is  formed,  filter,  wash, 
and  analyze  for  members  of  Group  II,  as  directed  on 
p.  94. 

Boil  off  all  traces  of  H2S  from  the  filtrate ;  test  for 
ferrous  iron^with  K3Fe{CN)g  ;  and,  if  it  be  present,  add 
a  few  drops  of  HNO3,  ^^^^  ^^^^  until  the  iron  is  wholly 
oxidized  to  the  ferric    state.     Unless  the    preliminary 


154  CHEMICAL  ANALYSIS 

examination  has  indicated  conclusively  whether  organic 
matter  or  phosphates  are  absent  or  present,  it  will  be 
necessary  to  test  for  them  at  this  point,  as  is  directed 
on  p.  109.  According  as  they  are  absent  or  present, 
follow  the  instructions  given  for  the  precipitation  and 
separation  of  the  members  of  Group  III,  on  pp.  110 
and  111. 

Boil  off  all  excess  of  NH^OH  from  the  filtrate  ^  which 
is  to  be  examined  for  Groups  IV,  V,  and  VI ;  add 
(NH4)2S  in  moderate  excess,  and  if  a  precipitate  is 
formed,  filter  and  wash  thoroughly.  Examine  it  for 
members  of  Group  IV  according  to  the  directions 
given  on  p.  118. 

To  the  filtrate  which  may  contain  Groups  V  and  VI 
add  NH4CI,  NH4OH,  and  (NH4)2C03  in  quantity  suffi- 
cient to  precipitate  completely  any  members  of  Group  V 
which  may  be  present.  Warm  the  mixture  gently  ;  and 
if  a  precipitate  has  formed,  filter  and  wash  with  ammo- 
niated  water,  rejecting  the  washings.  Examine  it  for 
members  of  Group  V,  according  to  the  directions  given 
on  p.  123. 

Concentrate  the  filtrate  which  is  to  be  examined  for 
Group  VI,  and  add  small  amounts  of  (NH4)2S04  and 
(NH4)2C204,  to  remove  any  traces  of  Ca  and  Ba  which 
may  be  present.  Filter  and  reject  the  precipitate,  if 
one  be  formed ;  and  examine  the  filtrate  for  members 
of  Group  VI,  as  directed  on  p.  127. 

Concentrate  some  of  the  original  solution,  and  test 
for  ammonium  salts  with  NaOH. 


SYSTEMATIC  PROCEDURE   OF  ANALYSIS       155 

SYSTEMATIC   ANALYSIS  FOR   ACIDS 

It  will  have  been  observed  that  the  preliminary  exami- 
nation and  the  results  of  the  analysis  for  metals  throw 
much  light  upon  the  nature  of  the  acids  which  may  be 
present  in  the  material  which  is  being  analyzed. 

For  example,  the  presence  of  tartaric  acid  may  be 
indicated  by  the  result  of  heating  in  a  closed  tube; 
nitrates  or  chlorates  show  their  presence  by  deflagration, 
when  heated  on  charcoal;  carbonates,  sulphites,  sul- 
phides, and  cyanides  are  detected  upon  the  addition  of 
HCl,  the  reagent  for  the  metals  of  Group  I,  by  efferves- 
cence with  or  without  characteristic  odor. 

Furthermore,  the  results  of  the  analysis  for  metals 
will  show,  according  as  Or  and  As  are  found  absent  or 
present,  whether  chromic,  arsenious,  and  arsenic  acids 
are  absent  or  possibly  present. 

But  in  addition  to  these  indications  there  are  others, 
depending  upon  the  nature  of  the  metals  present  in  a 
substance  and  upon  the  character  of  the  solution  of  that 
substance,  which  may  show  conclusively  whether  certain 
acids  or  groups  of  acids  are  absent  or  present.  If,  for 
example,  a  metal  of  Group  I  is  present  in  a  neutral 
or  acid  solution,  it  is  fair  to  presume  that  no  acid  of 
Group  II  can  be  present,  since  the  salts  of  Ag,  Pb,  and 
Hg^  with  such  acids  are  almost  universally  insoluble, 
either  in  water  or  acids.  If,  on  the  other  hand,  a  metal 
of  Group  V  be  found  present  in  a  neutral  solution,  it  is 
presumable  that  no  acid  of  Group  I  will  be  present, 
since  the  combinations  between  metals  of  Group  V  and 
acids  of  Group  I  are  all  practically  insoluble  in  water. 


156  CHEMICAL  ANALYSIS 

It  will  be  seen,  therefore,  that  a  knowledge  of  the 
solubilities  which  are  shown  in  Table  I,  p.  34,  will  save 
much  time  and  labor  by  aiding  in  the  interpretation  of 
the  results  of  the  analysis  for  metals  in  the  manner 
already  shown,  and  by  diminishing  the  number  of  acids 
for  which  individual  tests  must  be  made. 

In  proceeding  to  the  systematic  examination  for  acids 
it  is  desirable  to  remove  any  heavy  metals  which  may 
be  present,  since  they  are  liable  to  obscure  the  reactions 
expected  from  the  reagents  for  the  acids.  Accordingly, 
if  the  original  substance  is  soluble  or  partly  soluble, 
remove  the  heavy  metals  by  boiling  the  solution  with  a 
small  excess  of  NagCOg  and  filtering.  If  the  substance 
is  insoluble  or  partly  insoluble,  fuse  the  insoluble  mass 
with  Na2C03  in  a  platinum  foil  or  crucible,  boil  with 
water,  and  filter.  By  metathesis  all  the  heavy  metals 
become  carbonates,  and  the  alkali  metals  form  soluble 
salts  with  the  acids.  The  filtrate  from  either  method 
of  double  decomposition  can  now  be  analyzed  for 
acids. 

Fusion  decomposes  HgCgO^,  HgC^H^Og,  HCIO3,  and 
HC2H3O2 ;  but  salts  of  these  acids  are  soluble  in  water 
or  solvent  acids,  and  must  be  sought  for  in  the  portion 
soluble  without  the  aid  of  fusion. 

Of  course  it  is  necessary  to  test  for  HgCOg  in  the 
original  substance  before  NagCOg  is  added. 

Neutralize  a  small  portion  of  the  solution  of  the 
alkali  salts  with  dilute  HNO3,  and  heat  till  all  COg 
is  expelled.  Add  BaClg.  A  precipitate  indicates  the 
presence  of  members  of  Group  I.  Divide  a  larger 
portion  of  the  solution  into  three  parts,  and  test  for 


SYSTEMATIC  PROCEDURE   OF  ANALYSIS       157 

HgCrO^,  H2SO3,  and  H^SiO^,  as  directed  for  Sub- 
group I,  Group  I,  p.  138. 

If  any  members  of  Sub-group  I  are  present,  add 
dilute  HCl  to  another  portion  of  the  alkali  salts  solu- 
tion, and  pass  H2S  till  the  liquid  smells  of  it.  Boil  off 
excess  of  HCl  and  HgS,  and  divide  into  six  parts. 
Examine  for  H2SO4,  H3P04,H3B03,H2C204,  U^C^U^O^, 
and  HF,  as  directed  for  Sub-group  II,  Group  I,  p.  139. 

To  a  small  portion  of  the  alkali  salts  solution  add 
a  solution  of  ZnSO^  in  NaOH.  If  a  white  precipitate 
occurs,  treat  a  larger  portion  of  the  solution  in  like 
manner,  and  filter,  rejecting  the  white  residue.  Divide 
the  filtrate,  or  a  portion  of  the  alkali  salts  solution,  if 
H2S  is  absent,  into  three  parts;  and  analyze  for  HCl, 
HBr,  HI,  HCN,  H4Fe(CN)6,  H3Fe(CN)6,  and  HCNS,  as 
directed  for  Group  II,  p.  145. 

Divide  another  portion  of  the  alkali  salts  solution 
into  three  parts,  and  test  for  HNO3,  HCIO3,  HC2H3O2, 
as  directed  for  Group  III,  p.  150. 


158 


CHEMICAL  ANALYSIS 


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SEPARATION  OF  METALS 


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160 


CHEMICAL  ANALYSIS 


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ATOMIC    WEIGHTS 


161 


Table  IX  —  Atomic  Weights  of  the  Elements 


Aluminum 
Antimony 
Argon 
Arsenic    . 
Barium    . 
Bismuth  . 
Boron 
Bromine  . 
Cadmium 
Csesium  . 
Calcium  . 
Carbon    . 
Cerium    . 
Chlorine  . 
Chromium 
Cobalt     . 
Columbium 
Copper    . 
Erbium   . 
Fluorine  . 
Gadolinium 
Gallium  . 
Germanium 
Glucinum 
Gold  .     . 
Helium    . 
Hydrogen 
Indium    . 
Iodine 
Iridium   . 
Iron    .     . 
Lanthanum 
Lead  .     . 
Lithium  . 
Magnesium 
Manganese 
Mercury  . 


(F.W. 

H=l 

26.9 
119.5 

39.6 

74.45 
136.4 
206.5 

10.9 

79.35 
111.55 
131.9 

39.8 

11.9 
138.0 

35.18 

51.7 

58.55 

93.0 

63.1 
164.7 

18.9 
155.2 

69.5 

71.9 

9.0 

195.7 

3.93 

1.0 

113.1 

125.89 

191.7 

55.5 
137.6 
205.30 
6.97 

24.1 

54.6 
198,50 


CLARKE) 

Molybdenum 
Neodymium 
Nickel  .     . 
Nitrogen   . 
Osmium    . 
Oxygen     . 
Palladium 
Phosphorus 
Platinum  . 
Potassium 
Praseodymium 
Rhodium  . 
Rubidium 
Ruthenium 
Samarium 
Scandium 
Selenium  . 
Silicon 
Silver   .     . 
Sodium 
Strontium 
Sulphur     . 
Tantalum  . 
Tellurium . 
Terbium    . 
Thallium  . 
Thorium    . 
Thulium    . 
Tin  .     .     . 
Titanium  . 
Tungsten  . 
Uranium   . 
Vanadium 
Ytterbium 
Yttrium     . 
Zinc      .     . 
Zirconium 


H  =  l 

.  95.3 

.  142.5 

.  68.25 

.  13.93 

.  189.6 

.  15.88 

.  106.2 

.  30.75 

.  193.4 

.  38.82 

.  139.4 

.  102.2 

.  84.75 

.  100.9 

.  149.2 

.  43.8 

.  78.6 

.  28.2 

.  107.11 

.  22.88 

.  86.95 

.  31.83 

.  181.5 

.  126.1 

.  158.8 

.  202.61 

.  230.8 

.  169.4 

.  118.1 

.  47.8 

.  182.6 

.  237.8 

.  51.0 

.  171.9 

.  88.3 

.  64.9 

.  89.7 


162 


CHEMICAL   ANALYSIS 


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NOTES 


[The  heavy-face  and  light  numbers  below  correspond  to  pages  and 
reference  numbers  in  the  body  of  the  text.] 

9  1  Two  other  reasons  are  assigned  for  the  reaction  of  zinc  and 
sulphuric  acid  on  the  dikition  of  the  acid.  Dilution  increases  the 
ionization  of  the  acid  and  correspondingly  increases  its  solvent  action. 
Second,  impurities  in  the  zinc  prevent  polarization  and  enable  the  ions 
to  discharge. 

10  1  The  term  "colloid"  was  used  by  Graham  to  denote  substances 
"incapable  of  taking  crystalline  form,  and  also  distinguished  by  the 
mucilaginous  character  of  the  hydrates."  The  class  includes  glue, 
gums,  starch,  albumin,  and  certain  arsenic  and  cyanogen  compounds. 
These  compounds  are,  for  the  most  part,  of  great  molecular  weight. 
Crystalloids,  on  the  other  hand,  consist  of  easily  crystallizable  sub- 
stances, and  include  sugars,  salts,  and  strong  mineral  acids  and 
bases.  —  2  Ann.  d.  Chem.  u.  Pharm.,  121,  1.  — 3  A  porous  membrane 
consists  of  unsized  paper,  unglazed  porcelain,  or  parchment.  It  is 
pervious  to  most  solvents  and  to  crystalloid  solutes,  but  is  impervious 
to  colloids.  A  semi-permeable  membrane  is  either  a  colloid,  or  a 
porous  membrane  whose  pores  are  filled  with  a  colloid.  Semi-perme- 
able membranes  are  pervious  to  solvents,  but  impervious  to  both 
colloid  and  crystalloid  solutes. 

11  1  See  rfeffer,  Osmot  Untersuch.,  1887. — 2  Not  so  much  pro- 
toplasm itself,  but  rather  the  covering  of  the  protoplasm. 

12  1  Zeit.  f.  phys.  Chem.,  1,  481  (1887).  —2  Recent  data  by  Morse 
and  Frazer  from  experiments  with  cane  sugar  show  that  this  law 
should  be  somewhat  modified.  Amer.  Chem.  Journ.,  July,  1905,  — 3 
It  is  preferable  to  say  that  M  represents  the  gram-molecular  weight 
and  T  the  absolute  temperature. 

21  1  A  better  answer  to  the  question  is  that  the  velocity  of  the 
reaction  is  usually  increased,  and  frequently  the  nature  of  the  prod- 
ucts is  governed  by  the  temperature.  —  2  Inasmuch  as  a  substance 
in  solution  is  likened  to  a  gas  confined  in  a  closed  container,  dilution 

163 


164  CHEMICAL  ANALYSIS 

or  increase  of  the  volume  of  the  solvent  is  equivalent  to  increasing 
the  container,  or  decreasing  the  pressure. 

22  1  Arrhenius.,  Zeit.  f.  Chem.,  1  (1890).  —2  The  basic  character 
also  of  NH4OII  is  lowered.  In  a  few  cases,  however,  —  notably  in  that 
of  cobalt  and  nickel, — the  solvent  action  is  increased.  —  3  Dr.  Black 
suggests  a  better  explanation  of  this  reaction  :  "  HCl  and  NaC2H302 
are  both  highly  dissociated,  giving  high  concentrations  of  H+  and 
C2H302~  ions ;  but  H  •  C2H3O2  is  highly  dissociated,  or  H+  and 
C2H302~  ions  cannot  exist  in  great  numbers  in  the  same  solution. 
Hence  on  mixing  HCl  and  NaC2H302,  the  equation  Ci  •  C2  =  K  •  C3 
for  H+  and  C2H302~  ions  must  be  satisfied,  and  as  K  is  small,  C3  is 
large  and  Ci  •  C2  small.  By  increasing  C2  representing  C2H302~,  Ci 
representing  H+  may  be  still  farther  diminished ;  or,  the  more 
NaC2H302  is  added,  the  lower  the  concentration  of  H+  ions  and  the 
less  the  acidity  of  the  solution." 

28  1  Many  funnels  not  being  exactly  60°  will  not  fit  the  folded 
filter  described.  To  provide  for  angles  slightly  greater  or  smaller 
than  60°  in  folding  the  paper  the  second  time  to  form  a  quadrant,  the 
crease  is  made  so  that  one  half  the  fold  shall  be  shorter  than  the 
other.    This  can  be  made  to  fit  any  ordinary  funnel. 

45  1  As  suggested  on  p.  73,  these  tables  are  expressions  of  analysis 
by  the  dry  way.  They  are  not  necessarily  a  guide  to  exact  analysis, 
but  are  merely  for  preliminary  and  confirmatory  observation. 

47  1  The  best  way  to  separate  the  bead  from  the  wire  is  to  dissolve 
it  out  with  hot  water  acidified  with  hydrochloric  acid,  and  rinsed  with 
distilled  water.  The  loop  at  the  end  of  the  wire  should  be  permanent. 
Frequent  bending  of  the  wire  breaks  it.  When  coloring  the  bead,  if 
the  quantity  of  the  metallic  salt  or  oxide  is  large  enough  to  render 
it  opaque,  the  bead  can  be  "diluted"  by  taking  off  part  of  it  while 
hot  with  a  glass  rod,  and  building  it  up  again  with  more  borax. 

54  1  Compt.  rend.,  55,  576. 

56  1  Pogg.  Ann.,  119,  10. 

61  1  Mem.  II,  Phil.  Mag.  (4),  XXII,  329,  498.-2  Astronomy  and 
Astrophysics,  12,  321  (1893). 

66  1  Dr.  Montgomery  gives  the  following  instructive  definition  of 
equivalent  weight :  "Valence  is  power  to  combine.  If  two  amounts 
of  the  substances  are  equivalent  they  have  equal  power  to  combine 
or  do  chemical  work.  '  The  equivalent  weight  of  a  compound  is  that 
weight  which  interacts  with  the  equivalent  weight  of  an  element ' 
(Smith) .    As  the  equivalent  weight  of  an  element  is  that  weight  which 


'         NOTES  166 

will  combine  with  8  parts  by  weight  of  oxygen  or  with  1  part  by 
weight  of  hydrogen,  it  follows  that  the  equivalent  of  an  acid,  base,  or 
salt  is  the  molecular  weight  divided  by  the  total  valence  of  the  nega- 
tive, or  acid,  radical.  If  the  valence  of  the  radical  is  1,  the  com- 
pound will  interact  with  one  equivalent  of  an  element  or  compound  ; 
if  the  valence  is  2,  with  two  equivalents,  etc." 

68  1  Parsons'  Automatic  Gas  Generator,  originally  described  by 
Professor  Charles  L.  Parsons  in  an  article  entitled  "Distribution  of 
Hydrogen  Sulphide  to  Laboratory  Classes  "  (Journ.  Amer.  Chem.  Soc, 
p.  231  (1903)),  has  largely  solved  the  difficult  problem  of  delivering 
hydrogen  sulphide  to  classes,  when  used  in  connection  with  a  distri- 
bution system  easily  arranged  in  any  laboratory.  This  generator  is 
sold  by  Eimer  and  Amend,  New  York,  and  is  now  generally  used  in 
colleges  throughout  the  United  States.  The  apparatus  is  made  of 
stoneware,  and  a  single  generator  will  easily  supply  a  class  of  fifty 
students.  It  is  perfectly  automatic  in  its  action,  and  the  refuse  is 
automatically  removed.  The  pressure  on  the  hydrogen  sulphide  is 
almost  constant,  seldom  varying  more  than  two  or  three  millimeters. 
Only  fresh  acid  comes  in  contact  with  the  solid  reagent,  and  the  full 
strength  of  the  acid  is  utilized.  There  are  no  stopcocks  or  valves,  the 
apparatus  being  controlled  entirely  by  the  exits.  It  is  easily  cleaned 
and  refilled  without  taking  down.  It  requires  absolutely  no  attention 
for  weeks  at  a  time,  and  then  only  to  refill  the  acid  holder.  It  is 
always  ready,  and  generates  gas  only  when  in  use.  There  is  no  escape 
of  gas  while  recharging  the  acid. 

77  2  Throughout  the  remainder  of  the  text  all  mention  of  reagents 
refers  to  dilute  ones,  unless  otherwise  designated. 

78  1  Use  an  exceedingly  dilute  solution  of  ammonium  hydroxide 
(I  :  50).  At  first  the  precipitate  is  white  AgOH,  changing  quickly  to 
AgaO,  then  to  soluble  2  AgNOg  •  SNHg. 

79  1  Hg2S  is  formed  at  low  temperature,  but  is  broken  down  on 
warming.  —  2  The  precipitate  has  been  shown  to  consist  of  metallic 
mercury  and  white  Hg'sN  •  NO3.     J.  prakt.  Chem.^  39,  204. 

80  1  Pbl2  is  soluble  in  excess  of  KI.  Hence  the  precaution  of  add- 
ing reagents,  mentioned  on  p.  3  (/),  should  be  observed. 

81  1  Each  group  precipitate  should  be  thoroughly  washed  to  insure 
the  separation  of  that  group  from  remnants  of  the  filtrate,  which  might 
contain  members  of  subsequent  groups.  —  2  Ice  water  with  HCl.  Lead 
chloride  is  less  soluble  in  dilute  hydrochloric  acid  than  in  water. 
Reject  the  washings.  —  3  Instead  of  washing  out  the  lead  with  hot 


IGd  CSEMtCAL  ANALYSIS 

water  on  the  filter,  a  better  method  is  to  punch  a  hole  in  the  filter 
paper  with  a  glass  rod,  wash  the  contents  into  a  beaker,  add  water, 
boil  about  five  minutes,  and  filter  through  a  hot-water  funnel.  An 
ordinary  funnel  will  usually  serve  the  purpose.  —  4  In  the  presence  of 
much  Ilg'  and  a  little  Ag,  the  latter  may  not  be  found  here.  The 
residue  not  soluble  in  aqua  regia  should  be  examined  for  AgCl. 
Barnes,  Chemical  News,  51,  97  (1885). 

82  1  Hgl2  is  soluble  either  in  excess  of  KI  or  in  excess  of  HgCl2. 
Both  re-solutions  produce  double  salts,  Hg2Cl2l2  probably  being  pro- 
duced by  excess  of  HgCl2. 

84  1  This  compound  is  probably  CUSO4  •  4  NH3. 

87  1  KCNS  is  not  a  metallo-cyanide,  but  is  conveniently  classed 
here ;  and  besides,  with  ferric  iron,  it  forms  the  so-called  double  salt 
Fe(CNS)3  •  9  KCNS,  which  is  doubtless  a  metallo-cyanide. 

89  1  This  test  should  be  made  in  a  hood  with  a  good  draught  to 
insure  the  removal  of  the  poisonous  arsine. 

90  1  "  Colloidal  solutions  occupy  a  place  between  true  solutions  and 
mechanical  suspensions.  The  solute  is  present  as  particles  which  are 
so  extremely  small  that  they  can  neither  be  removed  by  sedimentation 
nor  by  ordinary  filtration.  By  warming  such  solutions,  and  by  addi- 
tion of  various  salts,  the  particles  may  be  made  to  increase  in  size 
until  they  are  precipitated."    Bailey  and  Cady's  Qualitative  Analysis. 

91  1  NH4MgAs04  is  soluble  in  acids,  and  hence  it  is  necessary  to 
neutralize  with  NH4OH. 

94  1  Experience  shows  that  many  of  the  sulphides  of  this  group  are 
soluble  in  warm  dilute  HNO3.  Hence  students  frequently  fail  to  pre- 
cipitate completely  all  the  sulphides  of  the  group  if  HNO3  is  not  com- 
pletely removed.  It  is  always  safe  to  evaporate  the  solution  to  small 
volume  to  expel  IINO3,  and  then  dilute  with  water  properly  acidified 
with  HCl.  — 2  The  quantity  of  hydrochloric  acid  in  this  connection  is 
important.  This  group  includes  eight  rather  widely  different  metals 
whose  sulphides  vary  both  in  their  solubility  and  in  their  chemical 
conduct  towards  reagents.  The  order  of  the  precipitation  of  their 
sulphides  from  cold  dilute  hydrochloric  acid  solution  is  that  of  the 
metals  mentioned :  lead,  cadmium,  tin,  bismuth,  antimony,  copper, 
mercury,  arsenic.  If,  however,  the  quantity  of  the  acid  is  increased 
and  the  solution  heated  to  nearly  boiling,  the  order  is  partly  changed. 
The  sulphides  of  arsenic  are  rendered  more  insoluble,  while  those  of 
the  other  metals  become  more  soluble.  In  certain  quantitative  deter- 
minations of  arsenic  its  separation  is  effected  by   the   use  of  hot 


NOTES  167 

concentrated  hydrochloric  acid.  As  stated  above,  the  arsenic  sul- 
phides are  colloidal  in  character  and,  as  is  the  habit  of  such  com- 
pounds, are  rendered  more  insoluble  by  heat  and  acids.  The  degree 
of  acidity  of  the  solution  should  be  about  1  of  concentrated  acid  to  10 
of  water,  for  the  precipitation  of  the  arsenic  sulphides,  and  an  equal 
volume  of  water  added  on  cooling  to  precipitate  the  other  sulphides. 
—  3  Should  the  solution  contain  nitric  acid,  ferric  salts,  chromates, 
and  other  oxidizing  agents,  a  white  precipitate  of  sulphur  would 
appear  here.  This  should  not  be  mistaken  for  the  sulphides  of  the 
group,  all  of  which  are  colored.  —  4  Dilute  a  portion  of  filtrate  (a) 
and  pass  in  H2S  again  to  make  sure  of  complete  removal  of  Group  II. 
The  filtrate  should  be  thoroughly  tested  for  remnants  of  the  group 
with  H2S.  Either  an  excess  of  hydrochloric  acid  or  an  insufficiency 
of  hydrogen  sulphide  might  prevent  the  complete  precipitation  of  the 
group. 

95  1  This  reagent  should  be  carefully  inspected  and  made  fresh  at 
frequent  intervals.  On  standing,  either  it  is  destroyed  by  oxidation, 
or  it  runs  to  the  higher  sulphides  by  the  loss  of  ammonia.  In  the 
former  case  it  becomes  colorless,  and  in  the  latter  case  the  yellow 
color  changes  to  red.  The  reagent  should  be  yellow. — 2  Though  CuS  is 
practically  insoluble  in  Na2Sx,  HgS  is  more  soluble  in  this  reagent  than 
in  (NH4)2Sx.  Hence  if  Hg''  is  suspected,  it  is  better  to  avoid  the  greater 
evil  and  use  (NH4)2Sx,  even  though  copper  may  be  present.  —  3  For 
detecting  traces  of  arsenic,  neither  the  mirror  nor  the  ammonium 
molybdate  test  is  very  reliable.  For  this  purpose,  Fleitman's  test, 
which  is  more  delicate,  is  given :  In  a  large  test-tube,  hydrogen  is 
generated  by  heating  a  concentrated  solution  of  sodium  hydroxide 
with  zinc  or  aluminum  (Gatehouse)  nearly  to  boiling.  Introduce  the 
arsenic  solution,  and  above  the  liquid  near  the  mouth  of  the  tube 
insert  a  loose  plug  of  absorbent  cotton  to  absorb  the  moisture. 
Spread  over  the  mouth  of  the  tube  a  cup  of  filter  paper  on  which  is 
placed  a  crystal  of  silver  nitrate.  On  warming  the  tube,  arsine  is 
evolved,  which  first  forms  a  yellow  coating  on  the  crystal,  quickly 
turning  black.  Another  delicate  test  is  Gutzeit's,  which  is  elaborately 
described  in  the  U.  S.  Pharmacopceia,  1900,  p.  521. 

96  1  Drs.  Emerson  and  Boggs  suggest  tin  for  zinc,  and  remark : 
"We  find  the  precipitation  of  antimony  by  means  of  metallic  tin  is 
better  than  zinc,  as  it  avoids  the  co-precipitation  of  tin  with  anti- 
mony."—  2  In  this  separation  the  zinc  must  not  be  allowed  to  dis- 
solve completely,  lest  the  tin  also  dissolve  and  oxidize  before  testing 


168  CHEMICAL  ANALYSIS 

with  HgCl2.    On  account  of  the  rapid  oxidation  of  SnCla,  it  should  be 
tested  quickly  with  HgClj,  after  filtering  and  dissolving  with  IICl. 

98  1  Only  a  small  quantity  of  nitric  acid  should  be  used,  lest  some 
mercury  be  dissolved.  Should  this  happen,  the  metallo-cyanide  of 
mercury  would  be  decomposed  by  hydrogen  sulphide,  and  thus  the 
test  for  cadmium  would  be  obscured  (cf.  p.  86). 

99  1  Residue  {a')  may  also  contain  sulphur  and  PbSO^. — 2  It  is 
preferable  to  add  H2SO4  before  evaporating,  which  would  drive  out 
HNO3.  — 3  Lead  should  be  entirely  precipitated  here,  lest  the  remnant 
again  appear  as  PbS  in  the  test  for  cadmium.  —  4  In  filtering  this 
mixture  care  should  be  taken  to  pour  on  the  liquid  slowly  with  a 
glass  rod  so  as  to  collect  all  of  the  precipitate  at  the  apex  of  the  filter. 
When  carefully  washed,  a  few  drops  of  boiling  dilute  hydrochloric 
acid  are  added  to  the  residue  and  filtered  into  a  beaker  with  about 
200  c.c.  of  water.  A  solution  of  sodium  chloride  hastens  the  precipita- 
tion.—  5  For  further  confirmation  of  the  presence  of  bismuth,  the 
white  precipitate  is  allowed  to  settle  and  is  filtered.  Then  a  solution 
of  sodium  stannite  is  poured  on  the  filter.  The  residue  turning  black 
confirms  bismuth.  Noyes  and  Bray  prescribe  the  following  prepara- 
tion of  sodium  stannite  solution  :  ' '  Add  a  10  per  cent  solution  of  NaOH 
solution  to  a  10  per  cent  SnCl2  solution  until  the  Sn(0H)2  first  formed 
is  dissolved.  The  solution  must  be  freshly  prepared."  —  G  Another 
method  for  the  detection  of  cadmium  in  the  presence  or  absence  of 
copper  is  as  follows  :  Whether  the  ammoniacal  solution  is  blue  or  not, 
add  hydrochloric  acid  just  to  acid  reaction,  then  some  iron  filings ; 
boil,  and  filter.  Test  the  filtrate  for  cadmium  by  passing  hydrogen  sul- 
phide gas.  —  7  Should  the  precipitate  be  black,  CdS  would  be  masked 
by  the  black  HgS  or  PbS.  In  this  event,  boil  the  precipitate  in  about 
25  c.c.  dilute  H2SO4,  add  50  c.c.  water,  and  pass  H2S.  A  yellow  pre- 
cipitate confirms  cadmium. 

101  1  For  further  discussion  of  this  subject,  see  Ostwald's  Founda- 
tions of  Analyt.  Chem.  and  Alexander  Smith's  General  Inorganic 
Chemistry. — 2  Drs.  Emerson  and  Boggs  comment:  "Boettger,  fol- 
lowing Ostwald,  attributes  the  solubility  of  Zn(0H)2  in  NH4OH  to  the 
formation  of  the  complex  ion  Zn(NH3)n  — n  varying  with  the  concen- 
tration of  ammonia.  The  solubility  of  Mn(0H)2  is  attributed  to  the 
same  cause  as  that  of  Mg(0H)2,  while  that  of  the  Co(OH)2  and 
Ni(0H)2  is  attributed  to  both  causes,  excess  of  NH4+  ions  especially 
tending  to  form  the  complex  ions.  The  explanation  of  the  decreased 
solubility  of  Al,  Fe"',  and  Cr  hydroxides  is  also  a  little  different; 


NOTES  169 

for  example,  when  A1(0H)3  dissolves  in  the  alkalies,  it  gives  the  ion 
AlOs".  Water  reacts  with  it  thus  :  AIO3  plus  3  H2O  give  A1(0H)3  plus 
3  (OH).  Therefore,  any  suppression  of  the  OII~  ions  (as  by  addition 
of  NH4CI,  which  is  a  product  of  the  reaction)  tends  to  cause  it  to  go 
to  completion,  with  the  formation  of  insoluble  A1(0H)3."  These 
writers  then  facetiously  conclude:  "It  looks  as  though  there  are 
enough  theories  for  everybody  to  be  suited." 

103  1  Theory  of  solution,  p.  18.  — 2  Lov^n,  Zeit.f.  anorgan.  Chem., 
2,  404  (1896). 

104  1  When  ammonium  hydroxide  is  first  added  in  excess  in  the 
cold  a  pink  solution  is  formed,  but  on  boiling  chromium  hydroxide  is 
reclaimed. 

107  1  Single  ferrous  salts  oxidize  so  quickly  in  the  air  that  it  is 
deemed  expedient  to  use  the  double  ammonium  ferrous  salt.  This, 
however,  has  the  disadvantage  of  introducing  an  ammonium  ion 
which  retards  the  precipitation  by  ammonium  hydroxide  and  so- 
dium hydroxide. 

108  1  Kriiss  and  Moraht  give  this  compound  as  Fe(CNS)3  •  9  KCNS. 
Ber.  d.  chem.  Ges.,  22,  206. — 2  Borates  and  fluorides  are  so  rarely 
encountered  in  ordinary  analysis  that  they  will  not  be  considered  here. 

110  1  Phosphoric  acid  is  tested  for  at  this  stage  in  order  to  deter- 
mine which  procedure  to  follow,  In  Absence  of  Phosphates  or  In 
Presence  of  Phosphates.  —  2  The  action  of  ammonium  chloride  with  * 
ammonium  hydroxide  on  the  solubility  of  certain  hydroxides  may  be 
better  understood  by  this  additional  note  to  the  discussion  of  the  sub- 
ject. Influence  of  Ammonium  Salts,  pp.  101-104.  Using  more  recent 
lettering,  let  Ci,  C2,  C3,  respectively,  represent  a,  6,  c  in  the  equation 
a-b  =  c  ■  k  (cf.  p.  18).  In  order  that  any  hydroxide  may  be  precipi- 
tated, the  product  C3  •  K  in  the  equation  Ci  •  C2  =  C3  •  K  must  reach 
its  maximum  or  so-called  solubility  product.  In  the  cases  of  various 
hydroxides,  averages  of  the  solubilities  of  classes  of  them  are  given 
for  comparison : 

1.  Na,  K 20  moles  a  liter 

2.  Ba,  Sr,  Ca 0.1  '^ 

3.  Mn,  Mg 0.0002  " 

4.  Fe,  Cr,  Al,  Zn,  Co,  Ki  .     .     .       0.00004  " 

Now  if  an  ammonium  salt  —  say,  ammonium  chloride  —  is  added  to  a 
weak  ammoniacal  mixture  of  these  classes,  what  will  happen  ?  Class  1 
is  highly  ionized  and  very  soluble ;  Class  2  is  less  ionized  and  less 


170  CHEMICAL  ANALYSIS 

soluble  than  Class  1 ;  Class  3  is  less  ionized  and  less  soluble  than 
Class  2 ;  Class  4  is  poorly  ionized  and  sparingly  soluble.  Classes  1, 
2,  and  3  are  all  better  ionized  than  ammonium  hydroxide,  and  when 
ammonium  chloride  is  added,  by  the  common  ion  principle,  the  excess 
of  Nn4+  ions  will  suppress  the  0H~  ions  from  the  better  dissociated 
hydroxides,  and  drive  them  back  from  the  solubility  product,  thus 
preventing  precipitation.  In  Class  4,  however,  the  members  are  so 
poorly  ionized  and  so  insoluble  in  water  that  Nn4+  ions  of  ammonium 
chloride  have  no  opportunity  to  suppress  the  0H~  ions  of  the  hydrox- 
ides. The  results  are  that  ammonium  chloride  in  the  presence  of 
slight  excess  of  ammonium  hydroxide  redissolves  manganese  and 
magnesium  hydroxides,  but  does  not  affect  those  of  ferric  iron,  chro- 
mium, aluminum,  zinc,  cobalt,  and  nickel.  All  of  these  conditions 
obtain  only  when  the  excess  of  ammonium  hydroxide  is  slight.  "When 
a  large  excess  of  ammonium  hydroxide  is  used,  the  conditions  and 
results  are  different  from  those  just  discussed.  Ferric  hydroxide  is 
practically  unaffected,  but  all  the  other  metallic  hydroxides  consid- 
ered are  more  or  less  redissolved.  The  re-solution  of  aluminum 
hydroxide  is  caused  by  its  amphoteric  or  dual  nature.  It  is  both  a 
weak  acid  and  a  weak  base,  and  dissociates  into  two  systems  of  ions. 
As  an  acid  it  loses  water  and  dissociates  into  H+  and  the  aluminate 
anion  AlOa";  as  a  base  it  dissociates  into  the  aluminum  cation  A1+ 
and  3  0H~.  Now  when  ammonium  hydroxide  is  added,  its  Oil"  ions 
unite  with  the  11+  ions  to  form  water.  But  the  system  C,,  •  Caio2  is 
constant,  and  the  suppression  of  its  H+  ions  causes  more  A1(0H)3  to 
be  dissociated  into  H+  and  A102~  until,  if  sufficient  ammonium  hydrox- 
ide is  added,  the  whole  of  the  A1(0H)3  runs  into  H+  and  A102~. 
Daring  the  progress  of  these  changes  the  free  NH4+  ions  unite  with 
AlOa"  ions  to  form  Cnh4  •  Caio2  of  high  solubility  product.  Should  a 
better  ionized  base  —  say,  sodium  hydroxide  —  be  used,  solution  will 
be  effected  more  quickly  and  with  less  of  the  reagent.  As  regards 
chromium,  zinc,  manganese,  cobalt,  nickel,  and  magnesium,  an  excess 
of  ammonium  hydroxide  unites  with  them  to  form  complex  metallo- 
ammonia  cations,  thus  requiring  more  molecules  of  the  hydroxides  of 
the  simple  cations  to  reach  concentration.  The  effect,  of  course,  is  to 
redissolve  these  hydroxides  in  varying  degrees,  nickel  being  affected 
most  and  chromium  least.  The  addition  of  ammonium  chloride  to 
an  ammonium  hydroxide  solution  of  the  hydroxides  in  question 
weakens  the  ionization  of  the  ammonium  hydroxide  by  suppressing 
the  0H~  ions.     This  prevents  the  re-solution  of  aluminum  hydroxide. 


NOTES  171 

On  the  other  hand,  the  increase  in  the  number  of  NH4+  ions  increases 
the  number  of  metallo-ammonia  ions  of  the  other  metals,  and  thus 
prevents  their  precipitation.  —  3  Barium  carbonate  is  added  to  insure 
a  complete  separation  of  the  iron  group  from  manganese.  In  the 
presence  of  ammonium  salts,  ammonium  hydroxide  does  not  pre- 
cipitate manganese  hydroxide,  but  the  tendency  of  the  latter  is  to 
oxidize  quickly  to  an  insoluble  basic  oxide,  which  frequently  precipi- 
tates it  out  of  due  course.  Barium  carbonate  forms  with  the  members 
of  the  iron  group  insoluble  basic  carbonates,  but  produces  no  precip- 
itate from  neutral  or  slightly  acid  solutions  of  manganese  salts.  — 
4  Filtrate  (6)  contains  barium  which  was  added  as  a  reagent,  and,  of 
course,  it  will  be  found  in  Group  V.  But  a  test  for  its  presence  was 
made  before  adding  BaCOs.  Phosphates  of  barium,  strontium,  cal- 
cium and  magnesium  are  insoluble  in  alkalies,  but  are  soluble  in  acids. 
Hence  if  these  salts  are  present,  they  are  precipitated  by  NH4OH  and 
redissolved  by  HCl.  For  this  reason  filtrate  (6)  should  be  combined 
with  filtrate  (a). 

114  1  Unless  care  is  taken  to  perform  the  experiment  in  the  cold, 
the  compound  ZnS04  •  4  NH3  will  be  formed  instead  of  (NH4)2Zn02. 
—  2  Green  and  more  stable  manganous  sulphide  may  be  formed  by 
boiling  the  pink  variety  with  an  excess  of  the  reagents,  ammonium 
sulphide  and  ammonium  hydroxide. 

118  1  This  washing  should  be  done  with  warm  water  containing  a 
little  (NH4)2S.  The  appearance  of  manganese  in  the  filtrate  is  due  to 
oxidation  or  the  formation  of  a  colloidal  precipitate  on  account  of 
having  no  "salt"  in  the  wash  water.  Also,  we  find  that  the  NiS  may 
appear  in  the  filtrate  where  (Nn4)oS  is  used  as  a  precipitant.  This 
may  be  precipitated  by  boiling  with  acetic  acid.  — 2  If  the  NaOH  is 
not  added  in  considerable  excess,  the  Zn(0H)2  will  reprecipitate  on 
boiling.  Hence,  according  to  Fresenius  it  is  best  not  to  boil,  but  to 
stir  with  cold  NaOII  solution. 

119  1  In  case  manganese  is  not  found  in  residue  (a),  this  brown 
precipitate  is  to  be  examined  for  manganese.  After  filtering,  if  the 
brown  color  persists,  it  is  probably  nickel  sulphide,  which  can  be  pre- 
cipitated by  boiling  with  acetic  acid.  —  2  Possibly  a  better  way  of 
confirming  manganese  is  to  evaporate  the  solution  to  dryness  and 
fuse  with  dry  Na2C03  and  KNO3  to  quiet  fusion  on  a  platinum  foil. 
A  green  mass  on  cooling  confirms  manganese. 

120  1  A  large  excess  of  KCN  should  be  avoided,  as  the  excess 
must  be  neutralized  with  NaBrO.    On  the  other  hand,  a  small  excess 


172  CHEMICAL  ANALYSIS 

must  be  added  both  to  precipitate  cobaltous  cyanide  and  to  redissolve 
it  as  a  double  cyanide. 

126  1  NII4OH  only  partly  precipitates  Mg(0H)2.  —  2  Sodium  co- 
baltic  nitrite,  Na3Co(N02)6,  may  be  substituted  for  HaPtCle  in  testing 
for  ammonium  and  potassium  compounds.  In  this  event,  substitute 
acetic  acid  for  alcohol. 

133  1  For  complete  descriptions  of  this  method,  see  Fresenius' 
Qaantitative  Cfiem.  Analysis,  ICth  ed.,  p.  117,  and  Crookes'  Select 
Methods  in  Chem.  Analysis,  p.  26. 

146  1  Edward  Hart  gives  the  following  excellent  method  for  test- 
ing chlorides,  bromides,  and  iodides  in  presence  of  each  other  :  The 
iodide  is  oxidized  with  ferric  sulphate  and  the  iodine  distilled  off  and 
tested  with  starch  paper  ;  the  bromide  is  then  oxidized  with  potassium 
permanganate,  and  the  bromine  removed  and  tested  by  distilling  in 
chloroform  ;  and  the  liquid  is  finally  tested  for  the  chloride  with 
silver  nitrate.    Amer.  Chem.  Journ.,  6,  346. 

1511  Professor  Slangier  adds  this  instructive  supplement  to  prelimi- 
nary tests  for  solids  :  "  If  the  substance  is  a  solid,  apply  the  follow- 
ing tests :  (1)  If  it  is  a  paint  or  pigment,  extract  for  several  hours 
with  ordinary  ether,  and  use  residue  for  subsequent  tests.  (2)  If  it 
contains  arsenic  or  other  volatile  constituents,  add  concentrated  sul- 
phuric acid  to  decompose  and  nitric  acid  repeatedly  to  oxidize,  heat- 
ing over  a  small  flame.  Continue  until  the  liquid  is  clear  and  colorless. 
Then  evaporate  until  fumes  of  SO3  appear.  Dilute,  filter,  and  proceed 
as  usual.  (3)  If  it  contains  aluminum  or  is  largely  mineral  matter,  as 
a  baking  powder,  add  concentrated  nitric  acid  and  heat  at  first  gently 
and  then  more  strongly.  Repeat  the  operation  several  times  until 
carbonaceous  matter  is  consumed.  Then  add  hydrochloric  acid  and 
boil.    Proceed  as  usual." 

152  1  Noyes  and  Bray  prefer  to  treat  the  residue  with  concentrated 
sulphuric  acid  and  hydrofluoric  acid  in  lieu  of  alkali  carbonates. 
In  some  respects  this  has  its  advantages,  but  the  method  requires  the 
frequent  use  of  expensive  platinum  crucibles  by  inexperienced  stu- 
dents ;  and  also  in  the  presence  of  barium  (and  to  some  extent  stron- 
tium and  lead)  difBcultly  soluble  sulphates  are  formed.  Journ.  Amer. 
Chem.  Soc,  29,  140. —2  In  the  event  the  two  solutions  produce  a 
precipitate  or  the  fusion  does  not  effect  solution,  the  aqueous  extract 
or  mixture  from  the  crucible  is  evaporated  to  dryness.  The  mass  is 
then  taken  up  with  boiling  water  and  filtered.  The  filtrate  contains 
the  chlorides  of  sodium  and  potassium,  and  those  of  barium  and 


NOTES  173 

strontium  may  be  present.  The  filtrate  can  be  tested  directly  for 
barium  and  strontium  (cf.  p.  124).  The  residue  may  be  silica  and  the 
chlorides  of  silver  and  lead.  Boil  with  ammonium  acetate  and  some 
dilute  acetic  acid.  Filter  and  wash  thoroughly  with  hot  water.  Test 
filtrate  for  lead  with  potassium  chromate.  Warm  residue  with  potas- 
sium cyanide  solution.  Filter,  and  test  filtrate  for  silver  by  adding 
excess  of  nitric  acid.  {This  must  be  done  in  a  hood  to  avoid  breathing 
the  deadly  prussic  acid.)  Transfer  the  residue  to  a  platinum  crucible, 
and  test  for  silica  by  adding  solution  of  hydrofluoric  acid  and  a  few 
drops  of  sulphuric  acid. 

153  1  In  order  to  guard  against  accidents  and  as  a  reserve  for  tests 
for  ammonium,  a  portion  of  this  original  solution  should  be  set  aside 
for  future  examination.  —  2  If  the  solution  is  alkaline,  it  may  contain 
various  solutes,  namely,  silver  salts,  sulpho-salts,  silicic  acid,  metallic 
hydroxides,  etc.  In  this  condition  the  solution  should  be  acidified 
with  nitric  acid,  and  if  a  precipitate  forms,  more  acid  should  be  added 
and  the  solution  warmed.  If  the  precipitate  does  not  redissolve,  the 
mixture  is  filtered  and  the  residue  examined  as  other  substances  insol- 
uble in  water.  The  filtrate  is  evaporated  to  expel  nitric  acid,  diluted 
with  water,  and  examined  as  other  acid  or  neutral  solutions.  —  3  No 
matter  what  the  valence  of  iron  was  originally,  at  this  stage  of  the 
analysis  it  must  necessarily  be  bivalent  or  ous.  See  reactions,  bottom 
of  p.  107. 

154  1  If,  on  standing,  a  brown  precipitate  is  formed,  filter,  and  test 
residue  for  manganese  ;  if  the  filtrate  remains  brown,  boil  with  acetic 
acid,  filter,  and  test  residue  for  nickel.     Use  filtrate  for  Group  IV. 


i 


INDEX 


PAGB 

Acid,  acetic 150 

boric 135 

carbonic 130 

chloric 149 

chromic 130 

hydrobromic  .  .  .  142 
hydrochloric  .  .  .  141 
hydrocyanic  .  .  .  143 
hydroferricyanic  .  .  144 
hydroferrocyanic  .  .  143 
hydrofluoric  .     .     .     .137 

hydriodic 142 

hydrosulphocyanic  .  144 
hydrosulphuric  .     .     .144 

nitric 148 

oxalic 136 

phosphoric     .     .     .     .135 

silicic 131 

sulphuric 134 

sulphurous     .     .     .     .134 
tartaric      .     .     .     .     .  136 
Acids,  analytical  classification 

of 75 

"     reactions  for .     .     .     .130 
"     analysis  for   .     .     .     .155 

Alcohol 72 

Aluminum 100 

Ammonium 126 

Ammonium  salts,  reagents  .  69 
Analytical  classifications  .  73 
Analytical  groups  .  .  .  .  75 
Antimony 91 


PAGE 

Aqua  regia 68 

Arsenic,  arsenic      ....  90 

"        arsenious  .     .     .     .  88 

Barium     '. 121 

"        chloride     ....  69 

Bismuth 83 

Blowpipe 40 

Borax .  72 

"     beads  ....       46,  105 

Bromine 70 

Cadmium 87 

Calcium     .     .     .     .  '  .     .     .123 

"        hydroxide      ...  70 

Carbon  disulphide  ....  72 

Chlorine 70 

Chromium      .     .     .     .     .     .  104 

"        salts,  reagents      .  71 

Cobalt 116 

"     nitrate 70 

Copper 84 

Crucibles 41 

Decantation 26 

Dissociation  .     .     .     .     .     .  13 

Ether 72 

Evaporation 38 

Filtration 27 

Flame 40 

Fusion 43 

Fusion  mixture 43 

Group  I,     acids      ....  130 

"II,       " 141 

"     III,      " 148 


176 


176 


CHEMICAL  ANALYSIS 


Group  I,     metals    .     . 

.     7^ 

"     II,        "        .     . 

.     82 

"     III,      "        .     . 

.   100 

"     IV,       "        .     . 

.  114 

"     V, 

.  121 

"     VI,       "        .     . 

.   126 

Ignition 

.     39 

Ions 

.     15 

Iron,  ferric    .... 

.  107 

"     ferrous       .     .     . 

.  107 

"     salts,  reagents    . 

.     70 

Lead 

.     .     80 

"    acetate      .     .     . 

.     70 

"    peroxide    .     .     . 

.     .     72 

Magnesium    .... 

.  126 

"         sulphate    . 

.     70 

Manganese     .... 

.     .  114 

"         peroxide    . 

.     .     72 

Marsh's  test  .... 

.     89 

Mercuric  chloride    .     . 

.     70 

Mercury,  mercuric  .     . 

.     82 

"         mercurous    . 

.     .     78 

Metals,  analytical  classif 

ica- 

tion  of     .    . 

.     .     76 

PAGE 

Metals,  reactions  for   .     .     .  77 
"       analysis  of.     .     .     .153 

Nickel 117 

Potassium 127 

"        hydroxide    ...  71 

"        salts,  reagents  .     .  71 

Precipitation 33 

Separations .26 

Silver 77 

"      nitrate 71 

Sodium 127 

"      hydroxide  .     .     .     .  71 

"      peroxide      ....  72 

"      salts,  reagents      .     .  71 

Solution 6 

Spectra 54 

Spectroscope 53 

Spectroscopy 52 

Strontium 122 

Tin  chloride 71 

"   stannic 93 

"   stannous 93 

Washing 29 

Zinc 114 


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